Download Praktikum in Allgemeiner Chemie für Biologen und Pharmazeuten

Practical course in General Chemistry for Students of Biology
and Pharmaceutical Sciences (ETH 529-1001-00 P)
1
Introduction........................................................................................................................................5
Properties of glassware..................................................................................................................9
Cleaning of glassware ....................................................................................................................9
Glass figuration............................................................................................................................10
Glass cutting ..............................................................................................................................10
Smoothing..................................................................................................................................10
Tube bending .............................................................................................................................11
Pipettes pulling ..........................................................................................................................11
Volumes of laboratory glassware ...............................................................................................11
Is volume a conserved quantity (do volumes add up)? ............................................................13
II Fractioning methods 1.................................................................................................................14
Precipitation of coagulates: Ni(OH)2, Cu(OH)2, Fe(OH)3, Cr(OH)3 .......................................14
Precipitation of calcium carbonate CaCO3, filtration with a porcelain Buchner funnel......15
Precipitation of AgOH, AgCl, Ag2CrO4 ....................................................................................17
Precipitation of colloidal (nanoparticulate) Prussian Blue......................................................17
Re-crystallisation of a mixture of KNO3 with Cu(NO3)2: purification ...................................18
Drying of CoCl22H2O or CuCl22H2O.....................................................................................18
III Crystalline solids and salts ........................................................................................................21
Solid mixtures – mixed crystals ..................................................................................................21
Mixed crystals ..............................................................................................................................22
Thermochromism of ((CH3CH2)2NH2)2[CuCl4] ........................................................................23
Solubility of NH4Cl, KNO3, solubility product of KClO4.........................................................24
Argentometric titration ...............................................................................................................25
End point indication of some argentometric titrations ............................................................27
Volhard method .......................................................................................................................27
Method of Mohr.......................................................................................................................27
Method of Fajans .....................................................................................................................27
Preparation of 0.05 M AgNO3 solution and its calibration......................................................27
Argentometric titration of Br-, I-, SCN- .....................................................................................28
Argentometric titration with instrumental detection...........................................................29
Titration with conductometric end point detection..............................................................29
IV Fractioning methods 2 ...............................................................................................................30
Condensation................................................................................................................................30
Distillation ....................................................................................................................................32
Distillation of a azeotropic two-component mixture ................................................................33
Vacuum distillation......................................................................................................................34
Sublimation ..................................................................................................................................35
V Volatile substances.......................................................................................................................36
Determination of melting and boiling points.............................................................................36
Preparation of volatile substances..............................................................................................38
Determination of the relative molar mass M of volatile substances by means of vapour
density ...........................................................................................................................................39
VI Acids – Bases...............................................................................................................................41
Acid and base definitions by Lewis and Brønsted-Lowry .......................................................41
Hydrogen chloride gas – concentrated hydrochloric acid .......................................................42
Hydrogen fluoride – hydrofluoric acid ......................................................................................43
Sulphuric acid ..............................................................................................................................43
Nitric acid .....................................................................................................................................43
Sodium hydroxide........................................................................................................................44
Calcium oxide, calcium carbonate .............................................................................................44
Ammonia ......................................................................................................................................44
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Properties of concentrated acids and bases...............................................................................45
NH3 ................................................................................................................................................47
CaCO3 ...........................................................................................................................................47
NH3/HCl........................................................................................................................................48
Lewis acids … Lewis bases..........................................................................................................48
Preparation of AlCl3. Reaction of AlCl3 with ether. Reaction of AlCl3 with KCl .................48
Proton transfer in aqueous solution...........................................................................................50
pK values, pH concept, strong acids and bases, weak acids and bases, multistage
deprotonation ...............................................................................................................................50
Strong acids ..................................................................................................................................51
Weak acids....................................................................................................................................51
pH concept....................................................................................................................................51
Solutions with fixed pH: buffers.................................................................................................52
pH measurement..........................................................................................................................52
Preparation of a phosphate buffer of pH = 7.30 and I = 0.16..................................................54
Preparation of potassium hydrogen tartrate ............................................................................55
Water as acid, water as base.......................................................................................................55
Acidimetric titration, titration curves, neutralisation curves..................................................57
VII Redox reactions.........................................................................................................................59
Redox reactions, solvent free or in aqueous solutions ..............................................................59
Thermal decomposition of potassium chlorate .........................................................................61
Preparation of CuCl ....................................................................................................................62
Preparation of potassium tetrachloroiodate(III) ......................................................................63
Redox reactions in qualitative analysis......................................................................................64
Disproportionation of H2O2, catalase.........................................................................................67
Standard reduction potential Fe(CN) 3- / Fe(CN) 4- ...............................................................68
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Permanganometric titration .......................................................................................................69
Permanganometric determination of oxalic acid, (COOH)2....................................................69
Iodometric titration (of Cu2+ solution).......................................................................................70
Coulometric titration, analysis of S2O32- and H+.......................................................................71
VII Ligand exchange and coordination chemistry .......................................................................73
Introductory experiments in coordination chemistry ..............................................................75
Preparative coordination chemistry ..........................................................................................80
Tetraammine nickel nitrite Ni(NH3)4(NO2)2 .............................................................................80
Potassium dioxalato cuprate(II) K2Cu(OOCCOO)22H2O .....................................................81
Metal indicators ...........................................................................................................................82
Dithizone as metal indicator .......................................................................................................83
Complexometric titration............................................................................................................83
Determination of the hardness of water by complexometric titration....................................84
IX Chromatography and liquid/liquid distribution .....................................................................86
Chromatography..........................................................................................................................86
Chromatographic separation of dyes.........................................................................................86
Liquid-liquid distribution ...........................................................................................................87
Determination of the distribution coefficient of iodine in the solvent system H2O/CH2Cl2 ..88
Ion exchangers .............................................................................................................................89
Ion exchange chromatography: separation of Cu2+, Ni2+, Fe3+................................................91
X Qualitative analysis......................................................................................................................93
XI Mineral synthesis and analysis................................................................................................100
Synthesis and analysis of struvite Mg(NH4)(PO4)6H2O .......................................................100
pk values of some acids at 25°C ..................................................................................................102
Standard reduction potentials ......................................................................................................104
3
Complex formation constants ......................................................................................................105
Solubility products.......................................................................................................................106
Conductivity data.........................................................................................................................107
4
Introduction
Objectives
The beginner’s course in inorganic and general chemistry is intended to teach the fundamental
methods of laboratory work to students of biology and pharmaceutical sciences and to make them
familiar with the important reaction types in inorganic chemistry. Restrictions in the availability of
instrumentation, laboratory space, personnel and time call for a certain degree of flexibility.
The course presents a variety of materials to the student. Quantitative analyses, partially carried out
with instruments, require accurate and clean working methods. Some physicochemical
determinations complete the scope of the course. The laboratory course and the first-semester
lecture have some common contents, albeit not very extensively. Laboratory work that is only
aimed at following procedures without asking for the chemical background is useless. In all
quantitative experiments the weights, instrumental results, calculations and observations must be
written down in a laboratory journal. Written exercises support the understanding of the theoretical
background and serve as a control.
Safety
Laboratory work can be dangerous, especially for the untrained beginner. The worst hazards
threaten the eyes. Therefore, it is indispensable to wear
SAFETY GLASSES
Splashes of strong bases often cause the loss of an eye. Only suction tubes and flasks, round flasks
and desiccators can be evacuated safely, all other vessel are prone to implosion. Glass splinters in
the eye are difficult to detect for the surgeon.
Many substances are poisonous. Heavy metal compounds like HgBr2, Pb(NO3)2 etc. are almost as
toxic as KCN. Solutions of toxic compounds are aspirated into pipettes with the aid of a balloon,
never orally. For work with gases and vapours like Br2, NO2, HCN etc. a ventilated hood must be
used, also with chlorinated solvents and benzene. Organic solvents are often flammable, and their
vapours, especially diethyl ether, can be ignited explosively by the flame of a Bunsen burner even
some meters away. Poisonous chemical waste is not to be disposed of into the sink but in special
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chemical waste containers present in each laboratory. Concentrated acids and bases, especially
H2SO4, should be diluted by pouring them slowly into an excess of cold water (never vice versa).
Heating Methods
It is often necessary to heat reaction mixtures in order to be able to observe phenomena or to
accelerate processes. This can be done by means of electric heaters (especially for flammable
solutions) or with a Bunsen burner. Large vessels (beakers, conical flasks) are heated on a support
equipped with a fireproof glass plate while test tubes can be exposed directly to the flame. In order
to avoid sudden eruptions of liquid during the heating of solutions in large vessels boiling aids must
be added. Test tubes are held in a wooden clamp and the flame is applied just below the liquid level
so that the formation of large bubbles at the bottom is avoided. Homogeneous heating is achieved
by continuous and gentle shaking. Electric heat sources should always be mounted such that they
can be removed quickly from the reaction vessel. An apparatus fixed on a stand has to be mounted
so high that the electric heater must stand on a socket (e.g. “Labor-Boy” in order to make contact.
In case of overheating the socket can be lowered or removed instantly.
The Bunsen burner is a versatile heat source. However, one has to know its properties in order to
use it efficiently. The air supply plays a crucial role. If it is completely closed, a bright yellow and
soot producing flame is obtained. It has the lowest temperature but should not be used for gentle
heating since it would spoil the equipment with soot. This air supply position, however, is most
suitable for the ignition. When the air supply is opened slightly the yellow emission disappears and
the flame changes to a homogeneous light blue. This state is useful for gentle heating. If the flame
produces still to much heat the gas supply can be throttled. When the air supply is fully opened the
flame appears to be two-component and is accompanied by a rushing sound. It has a blue core and
an almost invisible and very hot sheath. The hottest region is located a few millimetres above the
tip of the blue core and reaches about 1500 °C.
Quantities and Concentrations
In chemistry the properties of substances are most important, however, quantities and
concentrations are also crucial. Quantities are relevant for the amounts of conversion in reactions,
and concentrations determine reaction rates and equilibrium positions. Quantities describe absolute
numbers of atoms or molecules. Unfortunately balances do not provide these numbers, conversion
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factors are needed to obtain them from ordinary masses. The factors are called molar masses and
describe the mass of a defined number of atoms of an atom type. The defined number is called the
mol and it corresponds to a number of 6.023 • 1023 atoms. Atomic molar masses have the unit
g/mol, means they indicate the mass of a mol of a kind of atoms. Molar masses of molecules are
obtained by summing the individual atomic molar masses of the atoms in the molecule. In order to
determine the mass m of a certain number of mol n of a substance such that the material can be
weighed, we calculate m = n • Mg with Mg being the molar mass. If the number of mol of a known
mass of a substance has to be calculated, the expression
m
n=M
g
applies. Concentrations are usually given in mol per volume unit by chemists, because particle
reacts with particle, and not mass with mass. Concentrations are measures of density, they indicate
how frequently a kind of molecule is encountered per unit volume, and therefore its activity in
reactions. The higher the density, the faster the reactions. The typical concentration unit is mol/l,
abbreviated M. Unfortunately manufacturers of chemicals indicate concentrations in solution in
percent of weight (% wt.), together with the mass density of the solution in g/cm3. However, with
aid of the molar mass Mg the concentration in mol/l can be determined. Example: A solution of
hydrochloric acid, HCl in H2O, has the mass concentration cm = 36 % (wt.) and a density of
ρ = 1.19 g/cm3. The molar mass can be taken from a periodic table or similar. For HCl we find
Mg = 36.46 g/mol. Setup: because of the mass density one litre (1000 cm3) of the solution weighs
1190 g. 36 % of this mass are HCl, and this fraction, divided by the molar mass, is the number of
mol of HCl in 1190 g, which also corresponds to a litre. We obtain the number of mol per litre this
way, the molar concentration cn, and this is the desired answer.
cn =
ρ•1000 cm3 • cm
100 • Mg
In order to determine the required weight of a solid to make up a certain molar concentration in
solution, one proceeds as follows: the concentration and the necessary volume of the solution are
set by the experimenter, since the number of mol is n = c • V, with c in mol/l und V in litre. The
weight of n mol is m = n • Mg with n = c • V, summarised
m = c • V • Mg
Dilutions are calculated easily in this system. During dilution of a solution with concentration c1
and volume V1 by addition of solvent the number of mol does not change, only concentration and
volume do. Therefore n = c1 • V1 = c2 • V2 with c2 and V2 representing concentration and volume
after dilution. It follows
7
c2 =
c1 • V1
V2
It is recommended to practise this kind of calculations until they are carried out almost
unconsciously, since they are indispensable in normal laboratory work.
8
I Glassware
Types of glass
In the chemistry laboratory special glass types are in use. Since most of the glass has to be
thermally and chemically resistant borosilicate varieties (Pyrex, Duran etc.) are preferred. In order
to shape Pyrex or similar glass perfectly a natural gas/oxygen blowpipe is required for the more
complex work. Some simple figurations can be done with a Bunsen burner.
Abbreviations for glassware used in this manual (German version)
RG: test tube
BG: beaker
Properties of glassware
Evacuation should be applied only to suction flasks (thick-walled), desiccators, round flasks
and suction tubes. All other vessels like conical flasks, beakers, bottles, flat-bottomed flasks and
graduated flasks implode upon evacuation.
Glassware that can be heated: beakers, round flasks, conical flasks, suction tubes, porcelain
dishes. Only thin-walled test tubes withstand sharp temperature shocks up to 250 °C. It is not
recommended to heat thick-walled vessels like desiccators, suction flasks and mortars.
Volume indications on beakers and conical flasks are only approximate values. The precision of
graduated cylinders is sufficient for preparative work only. In quantitative analytical work volumes
are measured with graduated flasks (calibrated on filling).
Cleaning of glassware
Normally a household detergent is sufficient for cleaning. Afterwards the glass is rinsed with tap
water and finally with deionised water from a washing bottle. Never rinse directly under the
deionised water tap! This is waste of an expensive resource! The inner surface of burettes and
pipettes which are no more completely wetted are cleaned with ethanol or acetone.
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Rapid drying of moist glassware: consider first whether drying is needed at all! Normally this will
be the case for graduated pipettes only. These are dried simply by attaching the rear end to a
vacuum pump by means of tubing and the aspiration of a small piece of paper towel to the tip. This
prevents pollution by laboratory air sucked in. After about 5 minutes the pipette is completely
dried.
Glass figuration
The most simple glassblower work should be familiar to every student of the sciences. Pyrex glass
is simpler to figurate for the beginner than technical glass, even if a blowpipe is needed for more
complex work.
Glass cutting
Attention: during the breaking of glass the hands should be wrapped in a towel or a part of the
laboratory coat because of the danger of cuts by the broken edges! Glass has only moderate tensile
strength and from a surface cut it breaks easily upon a pull. Glass rods and tubing (up to 20 mm
diameter) are scratched with a glass cutter at the desired position (single scratch). The rod or tubing
is held with both hands such that the thumbs point at the surface cut and is pulled apart under slight
bending. Cut 4 pieces from glass rods with different diameters, length 15 to 25 cm, and two glass
tubes of 25 cm and one of 15 cm.
Smoothing
The broken ends are rather sharp-edged and must be smoothed. Rods and tubes are brought into the
flame slowly and sideways, at about a right angle, under continuous rotation.
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Tube bending
The 15 cm tube is heated on a long stretch at the centre in a large flame under rotation. Reduce the
rotation to a slight back and forth motion let the tube bend in the flame under its own weight. Never
force bending, instead increase heating.
Pipettes pulling
A narrow stretch of the 25 cm tube is heated vigorously under rotation. The tube is quickly
removed from the flame as soon as it becomes soft and flexible. It is immediately elongated by
pulling apart forcefully. After cooling it is broken at the centre and yields two pipettes the wide
ends of which must be smoothed finally.
Volumes of laboratory glassware
It is useful to know the volumes of common pieces of laboratory glassware which are frequently
needed. The measurement of these volumes helps to illustrate the precision of the indications
imprinted by the manufacturer. The calibration of apparatus for volumetric analysis should be
checked from time to time anyway.
Volume determination of test tubes and graduated cylinders
The tare (weight of the empty piece) of two test tubes of different size is determined with a
preparative balance by weighing them in a 250 ml beaker. The tubes are then filled to the brim with
deionised water and weighed again. The volumes are noted in the laboratory journal for later use.
Exercise: calculate the fill height for 5 ml of liquid in the tubes. Model: test tube = half sphere
(bottom part) + cylinder (body). The diameter is determined with a ruler.
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d = 2r
h
r
r
The tares of the graduated cylinders in the inventory are determined. They are filled with water to
the lowest numerical mark and weighed again. They are filled further to the highest mark and
weighed again. Determine the volumes and compare with the indications.
Exact volume of the 100 ml graduated flask
The tare of a dry 100 ml graduated flask is measured, then it is filled with deionised water to the
only mark and weighed. The contents are poured into a beaker, letting the flask after-drip for a
couple of seconds. The flask is weighed again. Determine the temperature of the water in the
beaker. Calculate the filled in and the drained volume with a precision of 0.01 ml.
Temperature ° C
Density of water (g/ml)
10
0.999700
15
0.999099
20
0.998203
25
0.997044
30
0.995646
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Exact volume of the 10 ml pipette
The tare of a small dry powder bottle with a cap is determined with the analytical balance. 10 ml
deionised water of known temperature are transferred into the bottle by means of the pipette.
Determine the total weight of the bottle and calculate the drained volume.
Is volume a conserved quantity (do volumes add up)?
The tare of a graduated flask is determined, together with its stopper. It is opened and 50 ml of
deionised water and 50 ml ethanol (96%) are filled in with the 25 ml pipette. The stopper is
plugged in again, held with a thumb, and the flask is tilted upside down and back about ten times to
ensure thorough mixing. The flask is set upright on the bench and the filling level is observed. Are
50 ml + 50 ml = 100 ml? Weigh the full flask and calculate the mass of the filling with the tare.
Determine the density of the mixture by taking a sample with the 25 ml pipette and draining it into
a beaker with known tare. Weigh and calculate the true volume of the mixture (density = mass per
volume).
The tare of a 100 ml graduated flask with its stopper is determined and then it is filled to mark with
water. 2 g sodium chloride (NaCl) are weighed into a small dry beaker. The salt is dumped into the
flask through a funnel which should not touch the water level. The flask is stoppered and the NaCl
is dissolved by repeated tilting of the flask. When the NaCl has dissolved completely the filling
level is measured. Determine the total mass and the density of the solution like before and calculate
the true volume. NaCl has a density of 2.165 g/cm3. Compare the theoretical volume /water +
NaCl) with the determined one. The volume deviation can be measured by taking the diameter of
the neck of the flask and the distance of the liquid level from the mark with a ruler Vcylinder =  r2
h).
How is a solution of exactly 100 ml volume prepared correctly from a weighed amount of
substance? Why do volumes not always add up?
13
II Fractioning methods 1
Precipitates, crystallisation, filtration, decantation, centrifugation, drying
Fractioning methods serve for the separation and purification of substances. The methods discussed
here are all related to solid-liquid separation: substances are precipitated, a compound is
crystallised, the liquid phase over a solid is decanted, a solid is centrifuged from a liquid are
operations which have to be carried out differently depending on the properties and the amount of
solid. Drying is another type of separation, the solvent as the liquid phase is evaporated.
Precipitation of coagulates: Ni(OH)2, Cu(OH)2, Fe(OH)3, Cr(OH)3
Filtration with a paper filter in the filtration funnel, filtration with folded filters
Folded filters serve for the quick filtration of small amounts of precipitates or impurities like lints
etc. from large volumes of liquid. Precipitate and filter are discarded.
400 ml of tap water are filled into a 600 ml beaker, 0.05 g of chromium(III)chloride CrCl36 H2O
are dissolved and a few drops of conc. NH3 are added to precipitate the Cr(III) as its hydroxide. The
hardly recognisable precipitate is filtered off with a 10 cm funnel and a folded filter.
Coagulates, which are non-crystalline precipitates, are filtered with a normal filter paper in a
funnel. The pore diameter of a filter paper is 10-2 to 10-3 mm.
For analytical purposes there exist special "ash free" filter papers.
150 ml tap water are filled into a wide neck conical flask and 0.05 g iron(III) chloride are dissolved.
The iron is precipitated as its hydroxide by addition of a few millilitres of 2M sodium hydroxide. A
round filter of 10 cm diameter is folded twice, fitted into a 6 cm glass funnel and moistened with a
few drops of deionised water such that it sticks to the glass. The iron hydroxide is filtered passing
the mixture along a glass rod such that the liquid level remains at least 5 mm below the edge of the
paper. Finally, the precipitate is washed cautiously with some deionised water.
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Precipitation of calcium carbonate CaCO3, filtration with a porcelain Buchner funnel
In a 100 ml beaker 2 g of calcium chloride CaCl22H2O are dissolved in 20 deionised water; in a
50 ml beaker 1.45 g sodium carbonate Na2CO310 H2O or 0.54 g water-free Na2CO3 are dissolved
in 20 ml deionised water. The contents of the 50 ml beaker are poured slowly under stirring into the
calcium chloride solution. The rubber seal is placed into the neck of a suitable suction bottle
(Buchner flask, thick-walled!) and the vacuum line is attached with thick-walled rubber tubing to
the connecting piece of the bottle. If the solution to be filtered contains corrosive or strongly
poisonous agents a second empty suction bottle or a gas washing bottle (scrubber) must be inserted
between suction bottle and vacuum line with a second piece of tubing. Foaming solutions must not
be allowed to enter the vacuum line because may inflict damage and high repair cost. A Buchner
funnel is inserted into the rubber seal and a fitting piece of filter paper is placed on the sieve. The
vacuum tap is opened causing the paper to be sucked onto the sieve. The contact is further
improved by splashing some water onto the filter which tightens the contact to the sieve. The slurry
of calcium carbonate CaCO3 is passed along a glass rod to the centre of the filter paper. The
remaining calcium carbonate in the beaker is shaken with some added water and also transferred to
the filter. This washing is repeated until almost no CaCO3 is left in the beaker. The filter cake on
the funnel is washed further by pouring 20-30 ml water over it. The vacuum is sustained for 5-10
min. more because it helps to dry the CaCO3. Finally, the vacuum tap is closed and the tubing is
pulled off the connecting piece while holding the Buchner funnel tightly because it is released
suddenly from the rubber seal upon vacuum breakdown! If the vacuum pump is of the wateroperated type (obsolete) the pump must not be turned off before the tube is removed from the
connecting piece since the water in the pump would slam back into the vacuum in the suction
bottle! The filter paper with the cake can be separated from the Buchner funnel by means of a
spatula.
15
Synthesis of crystalline PbI2, filtration with the glass frit of pore size "4"
1.6 g lead nitrate Pb(NO3)2 and 0.2 g potassium iodide KI are dissolved each in about 100 ml
deionised water in two 250 ml beakers. To the Pb(NO3)2 solution a drop of 2M nitric acid HNO3 is
added. Both solutions are heated simultaneously close to boiling and then united. Let the mixture
cool slowly and observe. The golden yellow crystals of PbI2 are collected on a glass frit with pore
size "4" and washed with a few millilitres of cold deionised water. The filtering method is in
principle the same as with the Buchner funnel. After removal of the product, the frit can be cleaned
by scratching with a spatula and passing hot (not warm) water through it from the bottom side.
Glass filter frit
Rubber seal
Vacuum
Vacuum
Suction flask
Cleaning
Filtration
For small amounts of coagulates or crystal the almost lossless centrifugation is more suitable than
filtration. The laboratory centrifuges can take up to 6 centrifugation tubes made of polymer (safe
against breaking, with screw cap).
Centrifuges must be handled carefully. They have to be installed upright on solid ground.
Rotating centrifuge parts store a lot of mechanical energy. Parts and splinters flying off a rotating
centrifuge have high velocities together with the corresponding kinetic energies. Therefore,
centrifuge loads must be balanced before they are spun. Always insert two tubes with the same
filling level into opposite holders. Centrifuges must be braked by hand only very gently, otherwise
the Coriolis forces will transfer rotational energy to the liquid and cause the precipitates to whirl up
again.
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Precipitation of AgOH, AgCl, Ag2CrO4
Decantation, centrifugation
In analytical chemistry, the precipitates formed in the following experiment serve also for the
qualitative identification of the ions involved since the reactions are characteristic.
A small test tube is filled to 0.5 cm height with 2M NaOH with a small pipette, and ten times more
0.2M AgNO3 is added dropwise under continuous shaking. The brown AgOH formed sediments
quickly. The overlying solution is decanted, deionised water is added and the mixture is shaken
again. The precipitate is allowed to sediment again, the solution is decanted. Finally the remains are
centrifuged and the rest of the solution is decanted completely from the compacted precipitate.
Selective precipitation and centrifugation of Cl- and CrO42- as silver salts
About 2 ml of the already prepared dissolved mixture of 0.1M sodium chloride NaCl and 0.1M
potassium chromate K2CrO4 are transferred into a small test tube and 0.2M silver nitrate solution
(AgNO3) is added dropwise, with shaking after the addition of each drop. White silver chloride
AgCl is formed initially. As soon as brown hue appears, the addition is stopped and the precipitate
is centrifuged. Without stirring up the AgCl, another millilitre of AgNO3 solution is added, which
precipitates now the chromate present as red silver chromate Ag2CrO4. The tube is centrifuged
again such that the chromate settles on top of the chloride. Solubility products of AgCl and
Ag2CrO4 are 10-10 M2 and 10-12 M3 respectively. Why is the numerical value for Ag2CrO4 smaller,
despite it is obviously more soluble than AgCl.?
Precipitation of colloidal (nanoparticulate) Prussian Blue
About 0.4 g potassium hexacyanoferrate(II) K4[Fe(CN)6] are dissolved in 100 ml water and a
solution of about 0.3 g iron(III) chloride FeCl3 6H2O in 50 ml water is added under stirring. The
deep blue material formed (a kind of ink) is so finely dispersed and not uniformly crystalline such
that it does not sediment and can hardly be centrifuged (try!). Suspensions of colloids with
diameters < 10-4 cm are called sols. They are not solutions. The presence of particles can be
recognised as follows: about 0.5 ml of the ink are transferred into a large test tube and diluted with
17
enough water such that the light of a microscope lamp or a white LED portable lamp can penetrate
the solution. When the lamp is pointed from the side at the tube in a dark environment (e.g. in a
ventilated hood) diffuse light can be seen shining out of the tube in the direction of the
experimenter. This is caused by the scattering of light at small particles (Tyndall effect). A dark
blue solution which is prepared from CuSO45H2O, water and NH3 solution does not show the
phenomenon. Large bright single reflections are caused by dust particles.
Re-crystallisation of a mixture of KNO3 with Cu(NO3)2: purification
Copper nitrate has a solubility of 244 g in 100 g H2O at 0 °C, potassium nitrate only 13.3 g in 100 g
H2O. The prepared mixture of KNO3 and Cu(NO3)2 6H2O contains considerably more KNO3 than
Cu(NO3)2 6H2O, the latter serves to simulate an impurity. Under appropriate conditions it is
possible to recover pure KNO3 by re-crystallisation, albeit with some loss.
To 10 g of the prepared copper nitrate/potassium nitrate mixture a few millilitres of water are added
in a large test tube. The mixture is heated gently and more water is added dropwise, until all salt
has dissolved. Now the solution is allowed to cool slowly, and colourless KNO3 crystallises.
Finally it is immersed into an ice/water bath to obtain as much KNO3 as possible. Consider how the
crystal are filtered best. Eventually, the whole operation has to be repeated.
A simple measure for the quality of the product is the qualitative detection of copper with ammonia
(NH3). 1 g of the re-crystallised KNO3 is dissolved in water and some drops of concentrated
ammonia solution are added. The stronger the blue colour, the more impure is the product.
Dry the moist KNO3 on a filter paper in air. Compare this process with the one in the following
experiment.
Drying of CoCl22H2O or CuCl22H2O
Drying can be interpreted as two somewhat different operations:
1) The removal of adherent solvent from moist freshly filtered crystals.
18
2) The removal of solvent molecules incorporated in crystal structures, e.g. crystal water.
In many cases adherent solvent can be removed simply by placing the filter cake on a dry piece of
filter paper where the solvent evaporates. With hygroscopic materials or for the removal of crystal
water, however, the water must be removed by heating, often assisted by vacuum. The heating is
necessary because of the evaporation enthalpy of the solvent, only small amounts of solvent can be
removed solely with evacuation. Large amounts of solvent must be adsorbed with a drying agent,
very large amounts are collected in a cold trap in frozen state (e.g. in food drying).
In order to remove the crystal water from one of the above substances, determine the tare of a 10 ml
round flask with a ground glass stopper (NS 14.5) on a preparative balance and weigh about 1.5 g
of the substance into it. The bent drying tube with ground joint is filled with water-free calcium
chloride as a drying agent, between two glass wool plugs. The ground joint is slightly greased and
plugged into the flask. This is immersed into a hot water bath (100 ml beaker, deionised water,
electric heater or burner with tetrapod). Vacuum is applied for about 1 hour with the membrane
pump (do not forget to add water to the bath periodically).
19
Extension clamp
Drying tube
Joint clamp
Calcium chloride
Rubber stopper with
drillling and glass tube
Glass wool
Vacuum
After completion the vacuum tube is removed, then the drying tube, finally the flask is sealed with
the ground glass stopper again. It is allowed to cool and weighed in order to determine the water
loss in the measures of mass and number of moles. The crystal water cannot be extracted from all
substances containing this. If the oxygen atom in the water is bound too strongly to a metal ion of
high charge the water is decomposed hydrolytically upon heating. An example is the reaction
AlCl36H2O

Al(OH)3 + 3H2O + 3HCl
Here the crystal water must be chemically decomposed, e.g. with thionyl chloride
H2O + SOCl2
2HCl + SO2
or the substance must be synthesised water-free with appropriate methods.
20
III Crystalline solids and salts
Appearance, mixed crystals, solubility, solubility product, enthalpy of dissolution
Crystalline solids are characterised by their regular structure. Such regular solids are known among
the metals, diamond-like materials (C, SiO2), the salts consisting of ions (NaCl, CuSO45H2O), the
refractory materials (CrCl3, CdI2) and also some molecular materials (I2, sugars). The regular
frame, the exactly repeating relative positions of the atoms, the so-called structure, can be
determined by X-ray diffraction analysis. This method yields size and form as well as the geometric
position of the atoms in the unit cell from which the whole crystal can be reconstructed. From the
superficial appearance of the crystal, the so-called habitus, only little information can be deduced
about the internal structure.
Solid mixtures – mixed crystals
Structurally equal particles of similar size often can substitute each other in crystal lattices. The
miscibility can be without limits. Examples are the systems potassium perchlorate – potassium
permanganate KClO4 – KMnO4 and potassium sulphate- potassium chromate K2SO4 – K2CrO4.
Even the charge of the substitute particle can be different as long as suitable compensating charges
are present. Miscibility is restricted here, however. An example is the substitution of Ba2+ by K+ in
BaSO4 under simultaneous substitution of SO42- by MnO4-. In contrast to this ordinary mixtures of
solids are separable, as shown in the previous chapter in the case of re-crystallisation.
A distinct limit of miscibility is found in the system (CuxZn(1-x))[Hg(SCN)4], with 0 < x < 1. Up to
the limit of x = 0.4 Cu2+ is incorporated instead of Zn2+ into the structure of the Zn[Hg(SCN)4].
There, the Cu2+ is in a tetraedric environment of four N atoms and has a deep purple colour.
If enough Cu2+ is added such that x > 0.4 the grass-green Cu[Hg(SCN)4], which has a totally
different geometry, is formed. Here the Cu2+ is in a square planar environment of four nitrogen
atoms and additionally bound to two sulphur atoms above and below of the CuN4 plane.
21
Mixed crystals
BaSO4 - KMnO4, CuxZn(1-x)[Hg(SCN)4]
In two large test tubes solutions of 0.1 g K2SO4 together with little (max. 10 mg) KMnO4 and of
about 50 mg BaCl22H2O, each in about 10 ml water, are prepared. Both solutions are heated
almost to boiling and the BaCl2 solution is added slowly to the mixed solution. The precipitate is
centrifuged. Try to wash the colour out of the pink precipitate by rinsing with water! This does not
work here since, in contrast to the previous recrystallisation of a KNO3/Cu(NO3)2 mixture, we have
a true compound. Let the crystals dry in air on a filter paper.
Preparation of mixed crystals of CuxZn(1-x) [Hg(SCN)4]
Prepare the following solutions in five test tubes:
1: 2.5 ml 0.1 M ZnSO4
2: 2.5 ml 0.1 M ZnSO4 + 1 drop of 0.1 M CuSO4
3: 2.0 ml 0.1 M ZnSO4 + 0.5 ml 0.1 M CuSO4
4: 0.5 ml 0.1 M ZnSO4 + 2.0 ml 0.1 M CuSO4
5: 2.5 ml 0.1 M CuSO4
To each of these test tubes 2.5 ml of 0.02M K2Hg(SCN)4 solution are added and the colours of the
precipitates formed are observed. The intensity of the purple colour caused by the direct
substitution of zinc by copper (x < 0.4) is remarkable.
22
Thermochromism of ((CH3CH2)2NH2)2[CuCl4]
Temperature changes can provoke structural changes in solids without the solid undergoing a phase
transition (e.g. melting). They are solid-solid phase transitions which do change physical properties
but not the composition. Rather spectacular are changes in light absorption. These are called
thermochromism. Several copper complex display this property caused by the four closest
neighbours (here: Cl-) of the Cu2+ ion switching from a square-planar environment to a tetrahedron,
thereby changing the light absorption properties.
Low temperature (green)
High temperature (yellow)
Preparation of ((CH3CH2)2NH2)2[CuCl4]: a slurry of 0.02 moles of diethylammonium chloride
(CH3CH2)2NH2Cl with 15 ml 2-propanol is prepared in a beaker. In a second beaker 0.01 moles of
water-free CuCl2 are mixed with 3 ml absolute ethanol. Both solids are dissolved completely by
gentle heating on a water bath. In the case of extended solvent loss this has to be replenished. The
diethylammoinum chloride solution is poured into the CuCl2. The mixture is allowed to cool to
room temperature and then immersed into an ice bath. If an oily precipitate is formed instead of
green crystals, the inside of the beaker under the liquid is scratched with a glass rod. The substance
often crystallises suddenly under this treatment. If thisdoes not work within 5 minutes, a few ml of
2-propanol are added and the solution is cooled again in the ice.
The green crystals are collected on filter paper in a Buchner funnel with vacuum assistance and
washed with ice-cold 2-propanol. Finally the crystals are dried in vacuum, transferred into a dry
glass vessel and dried overnight in a desiccator equipped with water-free CaCl2.
Demonstration of the thermochromic effect: 500 ml deionised water are filled into a 600 ml beaker.
A small test tube is filled with a clearly visible amount of the complex salt, this is compressed with
a glass rod and the tube is sealed with polymer film (Parafilm). The beaker is put on a tetrapod with
Ceran plate, the tube and a thermometer are mounted on stand with a clamp such that both a
immersed close to each other. The filling of the tube must be positioned complete below the water
surface! The beaker is heated now slowly with a Bunsen burner under occasional stirring with a
glass rod, and the complex is observed. Note the temperature when a colour change occurs. As
23
soon as the colour has switched completely, the burner is extinguished and the observation is
continued until the colour has switched back. Note also this temperature and compare the two
transition points.
Solubility of NH4Cl, KNO3, solubility product of KClO4
The solubility of a substance in a solvent can be formulated in different ways. The "Handbook of
Chemistry and Physics" tabulates the solubility as "grams per 100 g of solvent" for a given
temperature, since solubility depends strongly on temperature.
About 4 g NH4Cl are weighed into a large test tube and 10 ml water are added with a pipette. The
solid is dissolved by gentle heating and a thermometer is introduced. The solution is allowed to
cool now, and temperature is noted when the first crystals appear. Add two times1 g of NH4Cl,
repeat the dissolution and crystallisation after each addition, and note the crystallisation
temperatures for the different concentrations. In the case of KNO3 start with 4 g and add two times
3 g. Plot the solubility functions on finely graduated paper (mm grid).
Related with solubility is the enthalpy of dissolution, the solubility itself depends on the Gibbs
energy of the dissolution process.
For ionic solids with small to very small solubility the dissolution equilibrium constant is often
given instead of the solubility. The constant is called "solubility product", Kso.
MaXb(s)
a Mb+ + b Xa-
so: solubility
(s): solid
Kso = [Mb+]a [Xa-]b
[ ]: concentration in Mol/l (M)
This solubility measure is not only valid for the composition of the solution that results upon
dissolution of the salt but also for solutions with concentration ratios which differ from the
quotient of the stochiometric factors a/b. Kso is only active when the precipitate MaXb coexists
with the solution: if one concentration, e.g. [Mb+] is fixed, the constant also fixes [Xa-] by
precipitation or dissolution of MaXb. When saturation is reached upon dissolution or solvent
evaporation, solid exists in contact with dissolved: Solid and solution are in equilibrium, and the
position of this equilibrium is determined by Kso for a given temperature. In analogy other chemical
24
reactions like those of acids and bases, complex formations or redox processes always head into
equilibria, which are characterised by the corresponding laws of mass action with their constants
(compare with basic chemistry lecture). If the settling of the equilibrium is inhibited and lasts very
long, the reaction is said to be kinetically controlled.
0.48 to 0.52 g pulverised KClO4 are weighed exactly into a large test tube and dissolved
completely, beginning with 15 ml water, under heating. Cool to 25 °C. If KClO4 crystallises 0.5 to
1.0 ml water are added and the solid is dissolved again by heating. Repeat this procedure until a
saturated solution at 25 °C is obtained. The total weight of the water added is determined and the
solubility product of KClO4 is calculated.
Kso = [K+] [ClO4-]
for 25 ° C (density of the saturated solution d25 = 1.015 g/ml).
Argentometric titration
Application, end point detection and execution
Silver ions (also Hg22+ and Tl+) instantly form stochiometric crystalline and sparingly soluble
precipitates with many anions, e.g. Cl-, Br-, I-.
These properties are explored for quantitative analysis. Such sparingly soluble precipitates are
easily filtered, dried and weighed (gravimetry) or sample such as a halide can be titrated with silver
nitrate solution of known concentration with a burette (precipitation titration). The so-called
equivalence or end point has to be detected then. Several methods and effects are in use to
accomplish this. The following methods are non-instrumental:
For the determination of Ag+ by titration with a calibrated solution of KSCN (potassium
thiocyanate) some Fe3+ is added to the Ag+ sample solution. Towards the end of the AgSCN
precipitation the red iron(III) thiocyanato complex is formed (method of Volhard). For the titration
of Cl- and Br- with Ag+ a small amount of chromate is added as indicator. When the end point is
reached a red precipitate of Ag2CrO4 is formed (method of Mohr). The two dyes fluorescein and
eosin (both anionic) are also applied as end point indicators. At the passage of the end point the
25
surface of the silver halide crystals become charged positively by adsorption of excess Ag+ ions. In
turn the anionic fluorescein is adsorbed to the crystal surface and changes its colour due to this
phase transition (method of Fajans). A further method of end point detection is the so-called
clearing point. The AgX crystals formed initially in the titration are charged negatively because of
adsorption of excess X-. They repel each other and cannot coagulate (colloid). At the end point the
charge is neutralised and aggregates which sink rapidly can form. The fine colloidal turbidity
caused by the micro crystals vanishes. Reagent addition has to be carried out carefully and slowly
close to the end point, since a rapid overshoot only reverts the surface charge polarity and causes no
clearing (method of Liebig).
In the titration of cyanide CN- with Ag+ the colourless and soluble complex Ag(CN)2- is formed
initially.
2 CN- + Ag+
Ag(CN)2-
Upon further Ag+ addition, after half of an equivalent Ag+ per equivalent CN-, the precipitation of
AgCN(s) occurs.
Ag(CN)2- + Ag+
2 AgCN(s)
An example of instrumental end point detection is conductance measurement. During the
precipitation process the electric conductance remains more or less constant, since only ion
exchange takes place. After the end point, conductance rises linearly because of the addition of
excess non-reacting reagent.
Another detection method is based on the concentration dependent electric potential caused by a
silver electrode immersed in a solution of silver ions. This relation is expressed in the Peters
equation (logarithmic form of the famous Nernst equation):
RT
E = E°Ag+/Ag + nF ln[Ag+]
At the end point of an argentometric titration the free [Ag+] rises rapidly and causes a proportional
rise in electrode potential which can be used to indicate the end point. Potentiometric titrations do
not only allow for equivalence point determination but for the determination of solubility products
and complex formation constants, by evaluation of the complete titration function.
26
End point indication of some argentometric titrations
These indication reactions are also in use for the qualitative detection of the ions involved.
Volhard method
Some drops of 0.2 M silver nitrate and a crystal of iron(III) ammonium sulphate, Fe(NH4)(SO4)2
12H2O, are transferred into a large test tube and diluted with a few ml of water. Some crystals of
ammonium thiocyanate NH4SCN are dissolved in a medium test tube in a few ml of water. This
solution is added dropwise to the solution in the large test tube (shake after each addition) and
formation of the white precipitate of AgSCN is observed, together with the colour change at the
end point.
Method of Mohr
Here the "titration" is carried out the other way round: a solution of some crystals of KBr in some
ml of water is prepared in a large test tube, and 0.2 M AgNO3 is added dropwise. The indicator is a
1 M K2CrO4. What causes the brownish colour? Why is this compound formed only when all Brhas been precipitated?
Method of Fajans
Try the end point indication according to Fajans with KI. The experiment is set up like the method
of Mohr, except that 3 drops of eosine solution are added to the halide instead of K2CrO4 as an
indicator. Describe your visual impressions at the end point as precisely as possible in your
laboratory journal.
Preparation of 0.05 M AgNO3 solution and its calibration
Weigh about 2.2 g of AgNO3 into a large test tube, dissolve with some water and pour the solution
into a 250 ml graduated flask. Flush the test tube into the flask with some water. Fill the flask to the
mark, shake thoroughly and fill the 50 ml burette with the solution.
27
For the calibration about 300 mg KCl are weighed exactly (analytical balance) into a 100 ml
graduated flask, dissolved in water and filled to the mark. An aliquot of 25 ml of this solution is
transferred into a 200 ml conical flask with a pipette, diluted to about 150 ml with water and 1 M
K2CrO4 is added until the solution appears distinctly yellow.
A magnetic stirrer bar is placed into the conical flask which is set atop of a stirrer. About 15 ml of
the silver solution is added rather quickly under continuous stirring. Continue at a lower rate and
detect the end point according to the method of Mohr. Calculate the molarity of the silver solution
and compare with the concentration derived from originally weighed quantity and the flask volume.
Carry out two determinations.
Argentometric titration of Br-, I-, SCNCarry out an argentometric titration with the freshly calibrated silver solution. Samples are handed
out by the teaching assistant. Repeat until two results at least are in agreement. Method,
calculations and results are to be described in a short report.
0
10
20
30
40
50
5
4
6
5
7
4
3
8
3
2
9
2
1
11
6
7
8
9
1
10
Titration setup
28
Argentometric titration with instrumental detection
Argentometric titrations can be carried out with silver electrodes and millivolt meters if the number
of students per teaching assistant does not exceed 12, according to a separate manual handed out by
the teaching assistant, in small groups because of the limited number of instruments.
Titration with conductometric end point detection
Titrations with conductometric end point detection are no more carried out in basic courses since
acquisition and maintenance of the rather fragile and expensive (CHF 500.-) conductivity cells is
not affordable with the current student numbers.
The detection relies on the changes in electric conductivity of ionic solutions caused by reactions
which alter the concentrations of free ions. Try to imagine which changes the conductance will
undergo during an argentometric titration, given that contributions of all ion types are rather
similar.
29
IV Fractioning methods 2
Condensation, distillation, sublimation
The fractioning methods discussed in this chapter are based on the different volatilities of the
substances to be separated.
Condensation
Each one of the fractioning methods is used for different purposes. The condensation of a
component from a gaseous mixture at room temperature to the liquid state enables the use of this
liquid phase as a solvent. Further, water vapour can be condensed from gases by cooling, e.g. with
solid carbon dioxide/acetone, to almost complete removal. In a distillation the evaporated
substances must be re-condensed to liquid, which is rather difficult on the laboratory scale with
very volatile materials. In the next experiment gaseous NH3 is condensed as an example.
Condensation of NH3. Dissolution of Ca metal in NH3
Millions of tons of ammonia NH3 are produced every year in large industrial plants. The major part
is converted to fertilisers, polymers and explosives. Having a boiling point of -33 °C it is gaseous
under standard conditions. It is sold in steel cylinders liquefied under pressure (p = 300-400 kPa at
25 °C). Small amounts of ammonia are prepared in the laboratory by the reaction of ammonium
salts with OHNH4+ + OH-
NH3 + H2O
and condensation of the gaseous NH3 formed.
Liquid ammonia is a very good and in many respects water-like solvent for inorganic and organic
substances. Analogous to water it shows a slight autoprotolysis,
2 NH3
NH4+ + NH230
such that acid-base reactions are known in liquid NH3 (NH4+: strongest acid, NH2-: strongest base).
A solution of NH4NO3 in liquid ammonia behaves against non-noble metals like aqueous nitric
acid.
Rather remarkable is the dissolving capability of NH3 for alkali and earth alkali metals. They
dissolve in liquid ammonia with a deep blue colour without the formation of hydrogen. The blue
colour is caused by solvated electrons:
Ca
NH3
Ca2+(NH3) + 2e-(NH3)
The blue solutions are metastable and decomposed slowly under hydrogen evolution, faster under
catalysis. The reducing property of these solutions is revealed upon the addition of the acid NH4+,
one proton of which is reduced to molecular hydrogen:
2e- + 2NH4+
2 NH3 + H2
Ventilated hood !!
Assemble the apparatus shown below which is based on ground glass joints. The extension piece is
filled with KOH pellets sitting on a loosely stuffed glass wool ball, serving as a drying agent. The
receiving flask is cooled with solid CO2 (dry ice, -78 °C) which is ground well and mixed to a
slurry with 2-propanol (250 ml beaker, add 2-propanol to dry ice initially, not vice versa!). In order
to contact the NH3 gas to be condensed with the cold surface of the receptacle the introducing pipe
of the vacuum adapter must eventually be extended: attach a piece of PVC tubing. The 100 ml
recovery flask is filled with 15 g NaOH pellets and 10 g ammonium chloride NH4Cl. Mix the solids
rapidly by shaking, add some drops of water and attach immediately to the KOH drying tube. After
a while, about 2 ml of liquid NH3 are collected in the receptacle.
Liebig cooler
Solid KOH
Vacuum adapter
Extension piece
Glass wool
PVC tubing
NaOH + NH4Cl
2-Propanol
Dry ice
31
To the liquid NH3 a small piece of Ca metal is added. It dissolves slowly (cool receptacle
repeatedly in the cooling mixture) under the formation of a blue solution, in which blue-bronze
coloured phases can appear. Add some NH4NO3 with a small spatula. The blue colour vanishes
instantly under evolution of hydrogen, and remaining metal is dissolved completely:
Ca + 2 NH4+
NH3 Ca2+ + 2NH + H
3
2
At the end of the experiment the receptacle is placed in the opening of a ventilated hood and the
NH3 is allowed to evaporate.
Distillation
Distillation has two major applications. One is the purification of solvents, e.g. after the drying of
the solvent, thus avoiding difficult separation problems. Quite similar is the removal of the solvent
from a product after synthesis. On the other hand mixtures of volatile substances can be separated
by fractioned distillation (high reflux) even if their boiling points are close to each other. The ideal
case, where two miscible and volatile substances do not interfere in a distillation, is rare. In most
cases the mixture of two of these substances exhibits non-ideal boiling behaviour under formation
of so-called azeotropic mixtures which are characterised by a minimum or a maximum of boiling
temperature. The complete separation by distillation is impossible there. For example, all aqueous
solutions of hydrogen halides display azeotropic behaviour with a boiling point maximum when
distilled.
Temperature
Temperature
gaseous
Evaporation
Condensation
100% A
0% B
liquid
Boiling diagramm of an ideal
mixture
Azeotropic
ratio
0% A
100 % B
100% A
Boiling diagram of a mixture with
0% B
azeotropic boiling minimum
0% A
100 % B
32
Sparingly volatile substances can sometimes be distilled in vacuo, a method of which two variants
are known: distillation in a normal laboratory vacuum (p = 1.3 kPa) lowers the boiling temperature
by about 100 K, distillation in the vacuum of a rotary pump (p  0.2 Pa) affords a further boiling
point depression of about 30 K. According to the low pressures the gas volumes are expanded such
that vacuum distillations last longer than ordinary ones.
A special type of distillation can be used for separation and purification for sparingly volatile but at
least slightly hydrophilic substances. If the water in a – even heterogeneous – mixture H2O/X, with
X having only a small vapour pressure (e.g. nitrobenzene), is distilled off, the water molecules
carry molecules of X into the gas phase and to condensation in the distillate.
Such a steam distillation allows sometimes for the isolation of a snow white and pure product from
a kind of dirty tar.
The next experiment is an example for an ordinary distillation. The following paragraph describes
the execution of a vacuum distillation.
Distillation of a azeotropic two-component mixture
In this experiment an ethanol-water mixture, called wine, is used. The aim of the experiment is to
determine the azeotropic composition and the distillation temperature. The percentages of an
azeotrope depend on the distillation pressure. Here we distil under atmospheric pressure. The
composition of an azeotropic distillate can be determined from its index of refraction, a good
method, which requires a thermally controlled refractometer, however. Often the determination of
the density of the distillate is sufficient.
33
Thermometer
Liebig cooler
Vacuum adapter
Water bath
Cooling fluid
5
4
6
5
7
4
3
8
3
2
9
2
1
11
6
7
8
9
1
10
Receiving
flask
Set up the distillation apparatus based on ground glass joint according to the above drawing. A
water bath 400 ml beaker and Bunsen burner) serves as the heat source. Attention: ethanol is
flammable! About 50 ml of wine are transferred into the distillation flask, a boiling aid is added and
the winw is distilled. At the very beginning the ethanol-rich azeotrope is distilled off, water is
enriched in the distillation flask. As soon as about 2 ml of distillate are collected, the heat source is
removed, the receptacle is taken off and some of the distillate is sucked into a Pasteur pipette such
that no air is taken up from below. The liquid level in the pipette is marked with a felt tip pen kept
ready. The pipette content is expelled into a 50 ml beaker with known tare. The weight of the
distillate is determined immediately. The volume is determined by sucking water into the pipette to
the mark and weighing this amount of water like the distillate (assumption: (H2O) = 1 g ml-1).
Continue the distillation and note the temperature changes with time.
Vacuum distillation
For a classic vacuum distillation one needs a vacuum pump (membrane or rotary pump) and a
liquid trap, equipped with a manometer and a ventilation stopcock, between distillation apparatus
34
and pump. The joints of the distillation apparatus must be sealed with special vacuum grease. The
distillation apparatus should hold a capillary besides the thermometer. This serves to avoid sudden
eruptions by letting in fine air bubbles which are converted to vapour bubbles. Further possible
accessories are reflux condensers and distributing adapters in order to collect the fractions in
different receptacles without opening the apparatus. The usual heat source today is an electric oil
bath with controlled temperature, often combined with a magnetic bar stirrer. This kind of
apparatus is not available in the first semester course; most students will encounter it the first time
in the practical course in organic chemistry.
Sublimation
During sublimation a solid undergoes the transition into the gas phase directly, from which it recondenses as a solid. The process of evaporation and condensation is one-step. Sublimation is not
suitable for the separation of substances of similar volatility, but rather for the separation of nonvolatile components. Sublimation is often carried out in rotary pump vacuum. It is sometimes
difficult to remove the condensed solid from the vessel walls. Teflon tubes inserted into the
apparatus can be helpful. A simple sublimation is carried out in the acid-base chapter with AlCl3,
therefore we omit the experiment here.
A rule for the sublimation procedure is: never pass the vapour over a ground glass joint. Why?
Examples for substances undergoing sublimation:
HgCl2 (sublimate), NH4X, X = Cl, Br, I, iodine (I2), sulphur, benzoic acid, anthraquinone,
Al(CH3COCHCOCH3)3, camphor etc.
35
V Volatile substances
Melting point, boiling point, relative molar mass
Substances which can be transferred to the gaseous state at low temperatures consist of molecules
or atoms with only weak attractive interactions. These cause the coherence of the molecules in the
crystalline solid state. When the kinetic energy of the molecules is increased by rising the
temperature, the crystal lattice breaks at a certain temperature, the melting point. In the liquid state
the attractive forces still hold the molecules together, though they can slide over each other now. A
few molecules at the surface always acquire enough kinetic energy so that they can leave into the
gas phase. Their number grows with rising temperature until their partial pressure reaches the
atmospheric value, 101.3 kPa, and the liquid begins to boil. Melting and boiling point are
characteristic for molecular materials. If the boiling point is lower than the melting point the solid
phase is converted to the gaseous phase directly at 101.3 kPa: the material sublimes. Example: solid
CO2 (dry ice). The fact that also metals and salts have their (often high) melting and boiling points
does not contradict the above considerations. A far more characteristic value of molecular
substances is their relative molar mass. The knowledge of the molar mass helps to determine of the
molecular formula, e.g. Ne1, O2, P4, S2Cl2 etc.
Determination of melting and boiling points
In order to determine melting points two to three melting point capillaries are filled with substance
about 3 mm high. A filled capillary held by a clamp is immersed into a water bath (100 ml beaker,
see figure; stuff a piece of crumpled paper towel together with the capillary into the clamp to hold it
tightly). The water is heated slowly under stirring with a small gas flame. Use a thermometer
without ground glass joint. After melting of the first sample let cool the water by some degrees,
insert the next capillary and repeat the heating, rather slowly this time.
36
Thermometer
Melting tube
Water bath
5
4
6
5
7
4
3
8
3
2
9
2
1
11
6
7
8
9
1
10
Melting point determination
In order to determine a boiling point about 3 ml of the substance is transferred into a 10 ml
recovery flask, a boiling aid is added, a distillation adapter with ground joints equipped with a
ground joint thermometer is attached and the whole apparatus is fixed in a tilted position with a
clamp such that all condensing vapour flows back into the flask (see figure). Heat the liquid by
moving the Bunsen flame (weak setting) around the bottom of the flask until vapour condenses at
the thermometer tip. Boiling points, and to some small extent even melting points, are pressure
dependent. The observed values have to be indicated in the context of the air pressure.
37
Stand bar
Thermometer
Bosshead
Extension clamp
Liebig cooler
Joint clamp
Boiling point determination
Preparation of volatile substances
Structural characteristics as they can be found in crystalline solids are inexistent in ordinary volatile
substances. The most important method for the preparation of volatile substances is distillation
which has already been treated. On the other hand many experiments in this manual are concerned
with volatile substances such that a separate preparation of this type of material can be omitted.
38
Determination of the relative molar mass M of volatile substances by means of vapour density
"Molecular weights" of volatile substances can be determined, given almost ideal behaviour in the
gas phase, by the measurement of each of the quantities in the following equation:
pV = nRT
m
n=M
mRT
M = pV
The volume V of a glass container can easily be found by filling it with water to the brim and
weighing this filling. If molar mass determinations are carried out in an "open" container the
pressure is fixed by the ambient air pressure p, and if the "thermostat" is a boiling water bath the
absolute temperature T is also known. Only the mass m of the substance evaporated has to be
measured in order to calculate the relative molar mass M.
A 15 cm piece of thin copper wire, a 5 x 5 cm piece of aluminium sheet and a 200 ml narrownecked conical flask are placed on a preparative balance to determine the total tare. About 5 ml of
the substance under investigation is transferred into the flask which is sealed with the aluminium
sheet that is tied to the neck with the copper wire. The sheet is punctured at the centre with a fine
needle. The conical flask is clamped to a stand and immersed into a water bath (600 ml beaker) as
deeply as possible. During the following procedure it is indispensable to avoid water leaking
beneath the aluminium sheet. The bath is heated to boiling until all liquid in the flask has
evaporated. Leave the flask for 3 to 4 more minutes in the bath. This way the air almost completely
expelled from the flask. Now the flask is removed from the stand (clamp still attached to the flask)
and the flask is cooled under the water tap. The gaseous substance is re-condensed and the flask
refills itself with air. Let the flask come into thermal equilibrium with the environment for 2-3
minutes. Weigh the sum of flask, aluminium sheet, copper wire plus the mass of condensed
substance.
The flask is opened, blown out with dry air or nitrogen gas, weighed in empty state, filled with
water to the brim and weighed again to determine the total weight.
For the calculation the density of water can be assumed to be  = 1 g ml-1. The ambient air pressure
can be read from a barometer (e.g. at a switched off vacuum pump). The corresponding boiling
temperatures of water are:
39
p (kPa)
bp. H2O (K)
94.6
371.25
95.3
371.45
96.0
371.64
96.6
371.83
97.3
371.83
98.0
372.21
m(condensate) = m(flask+sheet+wire+condensate) - m(flask+sheet+ wire)
V(flask) =
M =
m(flask + water) - m(empty flask)
(water)
m(condensate)RT
pV(flask)
R = 8.314 J mol-1 K-1
(water)  1 g ml-1
Attention: the volume has to be converted to litre units for the molar mass calculation, and the mass
to kilograms! Despite the utter simplicity of the method the results are astonishingly accurate.
40
VI Acids – Bases
Acid and base definitions by Lewis and Brønsted-Lowry
Concentrated acids – concentrated bases
Lewis acids – Lewis bases
Proton transfer in aqueous solution, indicators
Acidimetric titration
Acidimetric titration, instrumental version
In chemistry historically there exist two definitions for acids and bases. The more recent and global
definition by Lewis names substances which can accept an electron pair from another substance as
Lewis acid. Examples: BF3, SnCl4, H+. Lewis bases can donate an electron pair. Examples: F-, NH3,
OH-. The concept is based on electron pair acceptors and electron pair donors.
The definition of Brønsted and Lowry is based on proton transfer: acids are proton donors, bases
are proton acceptors. Brønsted and Lewis bases are identical. The Brønsted acid concept is limited
to the single Lewis acid H+. From here on Brønsted acids are simply called acids. Two molecules
which differ only by one proton are called conjugate acid-base pairs, e.g. HCl – Cl-, H2O – OH-.
The acid-base definitions do include charges: acids and bases can bear positive, neutral or negative
charge.
The reactions between acids and bases, called neutralisations, for example
BF3 + F-
BF4-
(Lewis)
HCl + NH3
NH4+ Cl-
(Brønsted)
can occur solvent-free, e.g. in gas phase (HCl + NH3). In solvents it has to be distinguished whether
the solvent takes part in the neutralisation reaction or whether the solvent itself can be acidic or
41
basic, like water. Water can acquire a proton under formation of H3O+ (or its hydrated forms
respectively, abbreviated as H+) or lose a proton leaving OH- behind. The extent to which acids
transfer protons to water or bases extract them can be described quantitatively. This allows for the
classification of Brønsted acids (and bases) according to their strength.
We discuss the proton transfer in aqueous solution and the acid-base properties of water in theory.
We will apply acid-base reactions for analytical purposes in an acidimetric titration, and the
instrumentally assisted acidimetric titration will serve as an example of the quantitative
determination of acid strength.
The individual acids and bases show, especially in their concentrated forms, besides their proton
donating or accepting capabilities, some other chemical properties. These are listed below for the
most important concentrated acids and bases which are commercially as aqueous solutions, except
for sulphuric acid.
% Weight
Density Concentration(Mol l-1)
Hydrochloric acid (HCl)
36.5
1.19
12.0
Hydrofluoric acid (HF)
48
1.15
27.6
Nitric acid (HNO3)
65
1.40
14.5
Sulphuric acid (H2SO4)
98
1.84
18.0
Ammonia (NH3)
27
0.90
14.3
Attention: all these substances are very caustic. Wash splashes with much water immediately!
Wear safety glasses …
Hydrogen chloride gas – concentrated hydrochloric acid
Hydrogen chloride gas, produced from Cl2 and H2 or from NaCl and H2SO4, is sold in steel
cylinders. The aqueous solution which contains the ions H+ and Cl- is called hydrochloric acid. The
36% solution has a partial pressure p(HCl) = 101.3 kPa. Its chemical properties are based on 3
reaction types:
Protonation
Chloro complex formation
42
Redox reactions
Reduction of the proton to hydrogen
Oxidation of the chloride to chlorine
Hydrogen fluoride – hydrofluoric acid
Hydrogen fluoride (bp. 20 °C) is an extremely strong acid when water-free. The aqueous solution is
a weak acid. The ability to form complexes of the fluoride ion in hydrofluoric causes its glassetching property. The fluoride ion is not reducing.
Sulphuric acid
Stochiometric mixing in the ratio 1:1 of sulphur trioxide SO3 and water causes the formation of
sulphuric acid under considerable evolution of heat. The Lewis acid SO3 and the Lewis base H2O
combine to the protic acid H2SO4. The concentrated sulphuric acid has still a strong affinity
towards water. The rule: first water, then the acid … suits best to sulphuric acid. During the
distillation of pure H2SO4, SO3 escapes initially until the sulphuric acid is 98.3% (bp. 338 °C). This
concentration can also be reached by evaporation of water from dilute sulphuric acid (azeotropic
mixture, also found for the hydrogen halides). Sulphuric acid is a strong acid, the sulphate ion,
however, forms hardly any complexes. On the other hand, not only the proton is oxidising in
H2SO4, but also the hexavalent sulphur can, especially in the hot concentrated acid, act as a weak
oxidising agent and dissolve relatively noble metals like copper without hydrogen evolution.
Further, concentrated sulphuric acid has a strong affinity to water and is used as a catalyst in
esterifications or as a drying agent. The water is retained by protonation.
Nitric acid
100% nitric acid can be isolated. Distillation of aqueous solutions, however, yields only 69% HNO3
(bp. 122 °C), so-called concentrated nitric acid. With increasing concentration of nitric acid nitrate
ions and the protonated nitric acid molecule are formed by auto-protolysis.
2 HNO3
H2NO3+ + NO343
which decays to the nitronium ion and water:
H2NO3+
NO2+ + H2O
Since H2O occurs in the acid anyway there exists also the equilibrium
HNO3 + H2O
H3O+ + NO3-
The particle NO2+ is responsible for the strongly oxidising property of concentrated nitric acid.
With increasing dilution its concentration is lowered and the oxidising property vanishes.
Sodium hydroxide
Sodium hydroxide is sold as a solid in pellets. These pellets contain some water and cannot be
weighed exactly as 100% material. Solid NaOH vigorously attracts CO2 and H2O from the air. It is
soluble in water (42 g in 100 ml H2O at 0 °C) under enormous heat evolution (up to boiling!).
Concentrated solutions slowly attack glass. NaOH is a strong base, and OH- also acts as a
complexing agent. Never touch NaOH pellets with the skin. Always wear safety glasses when
working with solid or concentrated NaOH.
Calcium oxide, calcium carbonate
Calcium oxide CaO is a strong but sparingly soluble base and therefore hardly in use in the
laboratory. For large scale technical processes it is the cheapest base available. Often the even
cheaper calcium carbonate CaCO3, which contains the basic CO22- ion, is used to neutralise acid on
large scales.
Ammonia
Ammonia is sold usually as saturated aqueous solution. NH3 is a weak base and a good complexing
agent for many metal ions like Cu2+. Pure 100% ammonia is available in steel cylinders (bp. -35 °C,
101.3 kPa) or can be liberated from ammonium salts by strong bases:
NH4+ + OH-
NH3 + H2O
Liquid ammonia is a water-like solvent, though less acidic and with some interesting properties.
44
Properties of concentrated acids and bases
H2SO4: proton transfer to Cl-. Formation of HCl(g). Water-binding action: preparation of B(OCH3)3
In our first experiment the chloride ion in solid sodium chloride is protonated by concentrated
sulphuric acid under formation of gaseous hydrogen chloride. The hydrogen chloride is absorbed in
ethanol which serves as non-aqueous solvent, under formation of so-called ethanolic hydrochloric
acid. This used to produce free benzoic acid from sodium benzoate. The benzoic acid is reacted
with calcium carbonate in a later experiment. Pay attention to the conversion of the acid (proton):
O
H2SO4(l)
HCl(g)
H+Ethanol + Cl-Ethanol
+ CO32-
OH
H
CO2 + H2O
5 g of solid NaCl are transferred into the round flask of the apparatus shown in the drawing. The
dropping funnel is filled with 10 ml concentrated H2SO4 (ventilated hood). The receiving flask is
filled with 30 ml absolute ethanol. The sulphuric acid is allowed to drop slowly to the NaCl and the
gas evolution is finished by slight warming with the burner. The ground joint receiving flask is
sealed with a ground stopper.
45
vacuum adapter
PVC tubing
Bent joint
Ethanol
Dropping
funnel
Magnetic
stirrer
5
4
H2SO4
6
5
7
4
3
8
3
2
9
2
1
11
6
7
8
9
1
10
Ring for separatory funnel
Ground joints adapter
NaCl
Part of the ethanolic acid is used to convert sodium benzoate, the sodium salt of benzoic acid
O
O
Na+ + H+ + ClO
-
Ethanol
+ NaCl(s)
OH
thereby forming the ethanol-soluble benzoic acid while NaCl is insoluble in ethanol.
6 g of sodium benzoate are weighed into a conical flask, one half of the ethanolic hydrochloric acid
is added and the solution is shaken for some minutes. Filter off the insoluble residue. The filtered
46
solution is poured into a crystallisation dish and the ethanol is allowed to evaporate (in a ventilated
hood).
In the next experiment the water-absorbing property of sulphuric acid is used for the esterification
of boric acid with methanol (boron detection).
In a dry large test tube 10-100 mg of Borax Na2B4O710H2O (sodium tetraborate) are mixed with
1 ml concentrated sulphuric acid without heating, and 1 ml of methanol CH3OH is added slowly
drop wise and mixed. Add further 2 ml of methanol and dilute. Hold the test tube with a wooden
clamp, heat the mixture with the burner (ventilated hood) and ignite the vapours as soon as the
condensation level has reached the top of the test tube. The boric acid ester formed burns with a
beautiful green flame.
Borax
H2SO4
B(OH)3
CH3OH
B(OCH3)3
NH3
A volatile weak base. Hold the opened bottles of concentrated ammonia and concentrated
hydrochloric acid close to each other for a short time and observe.
NH3 as a complexing agent  chapter IX
NH3(l) as water-free solvent  chapter IV
CaCO3
CaCO3 is a very convenient starting material for the preparation of other calcium compounds from
free acids. Available in high purity degrees, stochiometric and not hygroscopic. The basic anion
CO32- is finally converted to the volatile CO2 and H2O by protonation.
3 g benzoic acid are weighed into a 100 ml conical flask, dissolved by addition of about 20 ml of
water and gentle heating. Add 5 g of CaCO3 divided in several quantities (weigh into a large test
47
tube). Finish the reaction by boiling shortly, filter the hot mixture through a small folded filter
paper and let cool slowly in order to crystallise Ca(benz)23 H2O. Filter and dry. Write down the
reaction equation.
NH3/HCl
NH3 and HCl in the water-free state react to, as seen before, the ionic solid NH4Cl(s), in water to
NH4+aq and Cl-aq. Upon heating of solid NH4Cl protons are increasingly transferred from NH4+ to
Cl- and the gaseous components NH3 and HCl are formed. At 340 °C the pressure of the gases
becomes so high that NH4Cl sublimes. In the same manner NH4Br and NH4I sublime at 452 and
551 °C, respectively. The increasing sublimation temperature is caused on one hand by the
increasing average molar mass of the gaseous components, on the other hand by the decreasing
affinity of Cl-, Br- and I- to bind protons: HI is the stronger acid than HBr, and this again stronger
than HCl, in the water-free state.
Transfer some 10-100 mg of NH4Cl into a medium-sized test tube, and sublime it in the full flame
of the Bunsen burner.
Lewis acids … Lewis bases
Preparation of AlCl3. Reaction of AlCl3 with ether. Reaction of AlCl3 with KCl
Aluminium trichloride AlCl3 can bind a ligand (Lewis base) with a free electron pair in a kind of
coordination expansion. AlCl3 is a Lewis acid. The Lewis base can be e.g. Cl-. The reaction with
NaCl
AlCl3 +
NaCl
NaAlCl4
subl. 183 °C
mp. 801 °C
mp. 152 °C
can be recognised easily from the change in melting and sublimation temperatures.
Ether CH3CH2OCH2CH3 can also serve as a Lewis base, of which one electron pair of the oxygen
becomes bound to the aluminium.
48
CH3
Cl
Cl
:
Cl
Cl
Al
:O
CH3
Al : O:
Cl
Cl
CH3
CH3
Lewis bases with higher affinity to aluminium can displace the already mentioned ones. The
aluminium etherate e.g. reacts vigorously with water under formation of the aquo complex
Al(H2O)63+. This one can be converted by fluoride F- into AlF63-, the hexafluoro aluminate ion.
This example demonstrates that complex formations are a special variant of Lewis acid-base
reactions (see chapter XI).
About 200 mg of crude water-free AlCl3 are transferred into a large test tube (close reagent bottle
tightly immediately after use). Add NaCl, about 1/10 of the amount of AlCl3. The opening of the
test tube is stuffed with a small ball of glass wool. Heat the bottom of the tube in the weak nonshining flame of the burner until AlCl3 sublimes to the upper cool part of the tube. The NaCl helps
to hold back impurities like FeCl3. AlCl3 may appear light yellow because of organic
contaminations. After cooling the sublimed AlCl3 is scratched out with a non-smoothed glass rod.
Half of the material is transferred into a medium sized test tube the tare of which was taken, and
weighed. Determine the equivalent amount of KCl, weigh it and add it to the AlCl3. Heat on a small
flame to formation of molten KAlCl4, mp. about 260 °C.
The rest of the sublimed AlCl3 is transferred into another test tube which is kept cool under the
water tap and 1-2 Pasteur pipettes of ether are added. AlCl3 dissolves under warming. The solution
is transferred into a small round flask by means of a Pasteur pipette. The flask is attached to the
vacuum pump where the excess of ether evaporates. The residue is solid aluminium chloride
etherate (mp. 36 °C).
The chloride in KAlCl4 as well as the ether in AlCl3(CH3CH2OCH2CH3) are substituted by H2O
molecules in a violent reaction when the substance is brought into water. This indicates that H2O is
the stronger Lewis base for Al3+ than Cl- or CH3CH2OCH2CH3.
49
Proton transfer in aqueous solution
pK values, pH concept, strong acids and bases, weak acids and bases, multistage
deprotonation
Acids are proton donors and are able to transfer one or more protons to a proton acceptor (a base).
For reasons of simplicity we shall discuss only acids with one proton initially; only one proton shall
be transferred. The relative tendency of the extent of a proton transfer can be measured versus a
standard base, e.g. against water.
Upon transfer of an acid HB into water the reaction (1) shifts into equilibrium:
HB + H2O
H3O+ + B-
(1)
The law of mass action for this reaction can be written as:
[H3O+][B-]
K' = [H O][HB]
2
Denominations:
HB = acid
B = corresponding (conjugate) base
H3O+ = hydrated proton, abbreviated H+
[ ] = concentration or activity in moles per litre = molar
K, K' = constants
If only dilute aqueous solutions are considered in which the concentrations of H3O+, HB and B are
smaller than 1 M the concentration of water in the solution can be regarded as constant.
[H2O] = constant
With K' [H2O] = K the law of mass action can be simplified to:
[H3O+][B-]
K =
[HB]
(2)
With the definitions pK = -log K and pH = -log[H3O+] the result, in logarithmic form is:
[B-]
pH = pK + lg [HB]
(3)
50
Strong acids
An acid is the stronger the higher its tendency to transfer protons is, the farther the equilibrium
position of reaction (1) lies to the right. In extremis reaction (1) runs almost completely to the right;
the corresponding acids are called strong acids. The strong acid hydrogen chloride HCl for example
forms, when brought into water, an equivalent amount of hydrated protons H3O+ and an equivalent
amount of chloride ions while HCl molecules are no more detectable. The K in equation (2)
becomes rather large therefore, and the pK value small: strong acids have small pK values
(pK < 0). The pH value of their aqueous solutions results, since for each HB one H3O+ is formed,
directly from the amount of acid added. pH = -lg[H3O+] = -lg[HB]added, with [HB]added being the
analytical concentration of the acid in the final liquid volume. The concentration is given in mol l-1.
Weak acids
If a weak acid is brought into water reaction (1) does not settle completely to the right. Only part of
the protons of the HB added is transferred to water. The particles H3O+, B- and HB are present in
similar concentrations, which are balanced under influence of the pK by equation (3). Die weak
acid dissolved in water forms less than the equivalent amount of H3O+ ions, the pH is greater than
the one of a solution of a strong acid at the same concentration.
A special case of a weak acid is water, which, according to the following equation
H2O + H2O
H3O+ + OH-
(4)
transfers a proton to itself. The law of mass action applied to this reaction yields
[H3O+][OH-] = Kw
Kw = 10-14 M2 at 20 °C, 101.3 kPa
pKw= -log Kw = 14
In pure water this process produces equal amounts of H3O+ ions and OH- ions.
[H3O+] [OH-] = [H3O+]2 = 10-14 M2
[H3O+] = 10-7 M
pH = 7
pH concept
The term pH = -lg[H+] is an important one in chemistry and technology. In chemistry it has an
influence on equilibria and kinetics, in biology it has characteristic values in body fluids, in
51
technology it is crucial in food preparation, enzymatic processes, and sewage treatment. Two
problems are to be solved in these fields: how to fix the pH in a solution to a known value, and how
can pH be measured?
Solutions with fixed pH: buffers
Solutions with very low (0 < pH < 3) or high (11 < pH < 14) pH values can be prepared from
strong acids and bases which are dissolved to the required concentration. In the range 3 < pH < 11
this method is useless: the amounts of acid or base become too small to compensate for changes
induced by further reagents. This is called an insufficient buffer capacity. Weak acids have the
advantage of partial dissociation and can provide low proton concentrations in the presence of
considerable total concentrations of acid. This is described by
[B]
pH = pK + lg [HB]
According to this equation the pH can be set for the given pK of the acid by adjusting the
concentration ratio of HB and B. Such mixtures are called buffers. Please note that the ratio
[B]/[HB] can be varied for a pH change of about pH = pK  1 without extensively losing buffer
capacity. It is clear from the expression that the addition of small amounts of a third acid (or a base)
will change the ration [B]/[HB] only insignificantly, the mixture will stabilise the pH, which is
called "buffering".
pH measurement
Two principles are frequently used, pH determination with the aid of electrode potentials (e.g. the
hydrogen electrode), especially the glass electrode, and pH determination with coloured acid-base
pairs, so-called pH indicators. If a small amount of such an acid-base pair HInd/Ind is added to a
solution containing the pair HB/B the ratio [B]/[HB] is hardly changed and therefore also the pH.
However, the pH of the solution determines the ratio of the added indicator components
[Ind]/[HInd] according to equation (3).
[Ind]
pH = pK + lg [HInd]
52
If Ind and HInd do have different colours the solution will acquire the mixed colour determined by
the pH. Inside the interval of about pH = pK  1 the pure colour of HInd changes through all mixed
colours to the colour of pure Ind. The hue can be used only for a coarse pH estimate by direct
visual observation. In this mode pH indicators are mainly used for the detection of pH jumps in
acidimetric titrations. Common indicators for the purpose are methyl red (red/yellow, pH = 5.0) or
phenolphthalein (colourless/purple, pK = 9.0).
With mixtures of several indicators of different pK values, so-called universal indicators, pH
estimates can be quickly found in the whole range in water.
Determination of the pK values of the indicator thymol blue
Thymol blue is an indicator which can release two protons stepwise:
H2Ind
pK1
red
H+ + HInd-
pK2
yellow
H+ + Ind2blue
For the determination of the pK1 four large test tubes are filled with 1 M, 0.1 M, 0.01 M and
0.001 M HCl, 10 ml each. Calculate the pH values. The dilute acid is prepared from 12 M
concentrated HCl. A fifth test tube is filled with 10 ml water. To each test tube an equal number of
drops of indicator solution are added. Estimate pK1 (trick: use a white sheet of paper as
background, and watch the colour from above).
In 6 large test tubes the following solutions are prepared:
Test tube 1: 10 ml deion. H2O
Test tube 2: ca. 100 mg NH4Cl
+ 1 drop NH3 conc. + 10 ml H2O
Test tube 3: ca. 50 mg NH4Cl
+ 5 drops NH3 conc. + 10 ml H2O
Test tube 4: ca. 20 mg NH4Cl
+ 10 drops NH3 conc. + 10 ml H2O
Test tube 5: 10 drops 2 M NaOH
+ 10 ml H2O
Test tube 6: 10 ml tap water
Add indicator solution and estimate pK2 of thymol blue (pK(NH4+) = 9.3). What is the approximate
pH of tap water?
53
Preparation of a phosphate buffer of pH = 7.30 and I = 0.16
Blood has an average pH value of 7.3 and only small deviations are tolerated. Furthermore, blood
contains ions, mainly Na+ and Cl-, producing an ionic strength I = 0.16 M.
1
I = 2 ciZi2
i
ci: concentration of the ith ion type in mol l-1
Zi: charge of the corresponding ion type
The second condition should be fulfilled e.g. in solutions for injection. Ionic strength, besides, is
the value needed to calculate activity coefficients. If an ion type is studied in solutions with ionic
mixtures it is found that the reacting ions behave as if their concentration were smaller than the
ones calculated from the weight. If the concentration of non-participating ions is further increased
the apparent concentrations of the reacting ions are decreased more and more. Explanation: ions in
solutions carry electric fields. These prevent that cations can approach each other, and the same is
true for anions. Ions of opposite charge can get closer than uncharged molecules normally do: they
form ion pairs and restrict the mobility of each other. A control of mobility exists, the statistical
probability to find an ion in a certain place is not equal for anions and cations and also different for
neutral molecules. Since concentration is a measure for the probability to find a molecule type in a
normalised part of space it can be no more representative for the reactivity contribution in an ionic
solution because of the mentioned local non-homogeneities. Ionic strength is a measure of the total
concentration of mobile charge in solution, and together with electric field theory a dimensionless
factor between 0 and 1 can be determined. It is multiplied with the concentrations. The factor is 1 in
dilute solutions (< 10-2 M) and approaches 0 for the highest concentrations of ions. For ionic
reactions one should always indicate activities for total concentrations above 10-2 M (= activity
coefficient  concentration). The concentration notation in square brackets we use (e.g. [Ca2+] is
actually reserved for activities.
Starting with disodium hydrogen phosphate Na2HPO412 H2O and sodium dihydrogen phosphate
NaH2PO42H2O 100 ml of a phosphate buffer of pH = 7.30 and I = 0.16 M are prepared.
[B]
1
Calculation: from pH = pK + lg [HB] the ratio [B]/[HB] can be obtained. With I = 2 ciZi2 we
i
can determine [B] and [HB] for the required ratio. Finally, with [B] and [HB] known, we can
54
calculate the needed amount of both phosphates. Measure the pH of the buffer prepared wit a glass
electrode and the pH meter.
Preparation of potassium hydrogen tartrate
The combination of free acid and fully deprotonated anion yields the monoprotonated anion in the
case of dibasic acids.
H2B + B2-
2 HB-
Examples are the formation of hydrogen sulphate or hydrogen carbonate:
H2SO4 + SO42-
2 HSO4-
CO2 + H2O + CO32-
2 HCO3-
and also the formation of hydrogen tartrate from tartaric acid and its dipotassium salt. Both are
highly water soluble while the monopotassium salt is sparingly soluble. Tartatric acid occurs to
have 3 structure isomers with different symmetry. The subject is more profoundly taught in the
organic chemistry classes.
Weigh 4.0 g tartaric acid and 6.3 g dipotassium tartrate in two 100 ml conical flasks each. Dissolve
both in a little water. Pour the solutions together, eventually through a small folded paper filter.
After a short interval the crystallisation of the monopotassium salt sets in. After some minutes to
allow for completion the product is filtered on a small Buchner funnel with filter paper, placed on a
suction tube. The precipitate is washed two times with a small amount of cold water and dried on a
filter paper in the air. From the monopotassium salt it is possible to obtain the mixed cations salt
KNa(tart)4 H2O, also called Seignette's salt.
Water as acid, water as base
The property of water to act as acid or base has already been mentioned. Strong acid donate their
protons completely to the base water. Besides of the protonic acids like H2SO4, HCl, H3PO4,
CH3COOH, H2tart, HSO4-, NH4+, H2S etc. there exist further classes of substances which generate
protons in their reactions with water. The non-metallic oxides like SO2, SO3, P4O10, CO2 etc. which
55
expand their coordination numbers under addition of the oxygen in water release protons thereby,
at a wide range of acidity.
SO3 + H2O
H+aq + HSO4-aq
H2SO4
The first proton released by the sulphuric acid formed is strongly acidic, the second bound to HSO4has to leave against the pull of the negative charge and is therefore less acidic. A solution of SO2 in
water is more a mixture of water and SO2 molecules, the solution is mainly molecular, not ionic.
Only the simultaneous active withdrawal of protons leads to coordination expansion:
-
-
SO2aq + H2O
+
HSO3 + H
OH
H2O
In the same manner water acts as a base against non-metallic halides and oxohalides:
PCl3 + 2 H2O
H3PO3 + 3 HCl
COCl2 + 2 H2O
CO2 + H2O + 2 HCl
Highly charged metal aquo ions, as they occur normally in aqueous solution, can liberate protons
from the water they bind:
Al(H2O)63+
Al(H2O)5(OH)2+ + H+
Under the influence of the positive charge of the central atom the protons of the water bound
become acidified.
The acidic action of water can be recognised in the reactions with basic anions:
O2- + H2O
CO32- + H2O
F- + H2O
2 OHOH- + HCO3OH- + HF
The solvent acts as donor of one proton here.
Transfer some milligrams or one drop of the following substances into a small test tube each, add
1 ml of water and one drop of universal indicator and note the estimated pH values.
NH4Cl, CaO, FeCl36H2O, K2C2O4H2O, CuBr2, CH3COONa3H2O, NaF, Na2S9H2O, Na2SO3,
NH3, Na2CO3, AlCl3
56
Write down the reactions with water which cause the observations.
Acidimetric titration, titration curves, neutralisation curves
If sodium hydroxide* is added in small quantities to the dilute aqueous solution of an acid HB the
pH increases after each dosage, and the concentrations [HB], [B] and [H2O] change because of the
neutralisation reactions:
H3O+ + OH-
2 H2O
HB + OH-
B- + H2O
If one plots as a result from such a titration the pH versus the volume of added base a titration curve
is obtained. For acids with pK values below about 9 the titration curves show a jump at the
equivalence point which can also be detected by the colour change of a suitable indicator.** This is
the foundation of a method for the quantitative determination of amounts of acid, since the number
of the moles of acid in the sample is equal to the number of moles of added base at the equivalence
point. Further the molar mass of an unknown acid can be measured, since the number of
equivalents can be calculated from the weight of the acid and its molar mass.
m(acid)
molar mass M of the acid = n(NaOH)
Titration curves can be displayed in a normalised form by using the neutralisation degree  as the
unit instead of the volume.
[OH-]add
 = [HB]
tot
At the beginning of the neutralisation curves obtained this way  = 0, at the equivalence point
 = 1. From the one-protonic acid handed out by the teaching assistant 1.5 – 2.0 g are weighed
exactly into 100 ml graduated flask, dissolved and filled to the mark. Estimate the pK from a small
sample of the solution by testing its pH. Evaluate a suitable pH indicator and titrate an aliquot of
10 ml of the acid solution in a wide-necked conical flask, diluted to about 150 ml with water, using
0.1 M calibrated NaOH in a 50 ml burette (determine twice). Calculate the molar mass of the acid.
In an analogue way, determine an unknown amount of acetic acid. Estimate the [H+] in vinegar.
* Use 0.1 M NaOH calibrated solution which is commercially available ("Titrisol").
** Later, compare the neutralisation curves recorded and determine the suitable pH indicator fort he corresponding titration.
57
The instrumental execution of acidimetric titrations with the aid of a glass electrode and a pH meter
allows for the quantitative recording of neutralisation curves. With this method it is possible,
besides the analytical application, namely detection of the end point jumps, to obtain data from
which exact pK values of unknown acids can be calculated, and equilibria coupled with
proronation/deprotonation reactions can be examined. These titrations are carried out in groups of
2-4 students. The experiments are described in a separate manual handed out by the teaching
assistant.
58
VII Redox reactions
Redox reactions, solvent free or in aqueous solutions
Electrode potentials
Permanganometric titration
Iodometric titration
Coulometric titration
Oxidation means an increase, reduction a decrease of the stochiometric valence of an atom. About
the assignment of oxidation numbers: see classes in chemistry. These processes are accompanied
by uptake or release of electrons, and since electrons cannot be set free except in a vacuum there
are always two half-reactions coupled. Redox processes occur in many ways: most of the methods
to generate chemically mechanical or thermal energy rely on redox reactions, e.g. cell respiration,
thermal power plants, combustion motors, heating facilities etc.
Some ions, atoms or molecules are known to take up or release electrons without other changes,
especially of the atomic composition. For example, Ce4+aq and Ce3+aq as well as Fe(CN)63- and
Fe(CN)64- are distinguished only by their contents of electrons, and in the redox reaction
Ce4+ + Fe(CN)64-
Ce3+ + Fe(CN)63-
which occurs in aqueous solution only electrons are transferred. The change in valence of an atom,
however, often causes changes in its close environment, so-called "coordinative rearrangements".
When the permanganate ion MnO4- is reduced to Mn2+aq the oxide ions O2- bound to the
managanese must react and be converted to water by protonation during the complete redox
process:
MnO4- + 8H+ + 5e-
Mn2+ + 4H2O
Here reduction depends on coordination changes which in turn cause acid-base reactions. Often the
single partial steps of such reactions are not really known, only the total reaction is. Compare this
59
reaction type with the synthesis of KICl4, the reaction of IO3- with I- and the permanganometric and
iodometric titrations in the following experiments.
Redox reactions are also known in heterogeneous systems. They occur at the surface of solids.
Molten lead chloride PbCl2 can oxidise metallic aluminium under formation of lead metal and
AlCl3, or Cu metal slurry can reduce Cu2+ to Cu(I) in chloride containing solutions. The oxidation
of coal with O2 to CO2 is classic heterogeneous redox reaction.
Another variant of reductions and oxidations results from the possibility to transfer electrons from
or into the surface of an electric conductor, namely an electrode. Electrons are released from or
taken up by the electrode surface. Making electrons available means reduction, taking up means
oxidation. The electrode material can be inert (means that it does not take part in the electrode
reaction) or it can be directly involved in the electrode reaction. As electrode reactions redox
processes can be spatially separated into the reduction and oxidation half-reactions.
Examples:
2H+ + 2e-(Pt)
Zn0
H2
reduction half cell at inert cathode
Zn2+ + 2e-
oxidation half cell with Zn0 anode
The electrolytic production of chemicals, e.g. ClO- or metal coatings (chromium plating),
electrogravimetry, energy release from batteries are applications of electrode processes. The
example of coulometric analysis show at which degree of refinement electrode processes can be
used for analysis.
Further uses of reactions at electrodes:
The potentials acquired by electrodes depend on the concentrations of the particles involved in the
electrode reaction. Potential measurements therefore allow for the determination of those
concentrations, sometimes to very small values. The determination of silver ion concentrations
[Ag+] with silver electrodes (instrumental argentometry) is such an application which enables,
besides the analytical determination of the end point, the determination of stochiometry,
equilibrium constants, solubility products etc. of reactions involving Ag+.
Electrode potentials measured or calculated under standard conditions (25 °C, c = 1 M, p = 101.3
kPa) can be used to assign a "standard reduction potential" to each reductive half-reaction. With the
arbitrary assignment of the zero value to the reaction
60
2H+ + 2e-(Pt)
H2
E0  0 V
and the positive direction to more oxidising systems the half-reactions can be listed according to
their reducing or oxidising power. In one of the following experiments the standard reduction
potential of the half-reaction
Fe(CN)63- + e-
Fe(CN)64- is determined.
Thermal decomposition of potassium chlorate
Potassium chlorate, a strong oxidising agent, is decomposed above its melting point (386 °C),
eventually under formation of KClO4, to O2 and KCl, which melts at 776 °C.
In order to formulate the reaction equation it is recommended to proceed as follows: the oxidation
states in starting materials and products have to be determined initially. The potassium ion K+
obviously does not take part in the reaction and is not considered further. Chlorate is ClO3- ; oxygen
is more electronegative than chlorine and obtains its lowest oxidation number, -II, which also
makes it fulfil the octet rule. Since the ion has a total charge of -1, chlorine is assigned an oxidation
number of +5. The products are chloride Cl- and dioxygen O2. The oxidation number of chlorine is
-1 therefore, and oxygen becomes elemental, has oxidation number 0.
We can write provisionally:
Cl(+V)O(-II)3-
Cl(-I)- + O(0)2
In order to keep stochiometry correct, we separate oxidation and reduction processes. This step
bears no relation to reality and is purely formal.
Reduction:
ClO3- + 6e-
Cl- + 3O2-
We pretend as if nothing would happen to the oxygen in the first place, only chlorine is reduced, it
takes up electrons.
Oxidation:
2O2-
O2 + 4e-
We assume that the oxygen in chlorate had completely taken over the valence electrons of chlorine
and it could dissociate as O2-. This state we oxidise formally to O2, the real product.
Now we recognise that the two partial reactions are not stochiometrically equivalent yet, because
the reduction requires 6 electrons, while the oxidation yields 4 of them only. In order to write the
total reaction, we have to balance the number of electrons, since “free” electrons do not occur in
61
normal chemical reactions. The least common multiple of 4 and 6 is 12. We have to multiply the
reduction equation with 2 and the oxidation equation with 3:
2ClO3- + 12e-
2Cl- + 6O2-
6O2-
3O2 + 12e-
These two equations we can add, almost as two algebraic equations:
2ClO3- + 12e- + 6O2-
2Cl- + 6O2- + 3O2 + 12e-
Identical ions or molecules on both sides can be subtracted according to the lower number. The
number of electrons should be the same on both sides because of the stochiometric adjustment and
therefore vanish. Accidentally, the numbers of O2- ions is also identical on both sides and O2- does
not appear in the final total equation. This has a meaning for the real reaction: if O2- were required
on the starting material side, the reaction would not start without the addition of extra O2-, e.g. in
the form of an oxide. The reaction, however, runs on KClO3 alone, as required by the reaction
equation. The final form is:
2ClO3-
2Cl- + 3O2
and describes very simply the observed decomposition. The real elemental steps of the reaction,
which can be very complex, are not captured this way.
Tare a small test tube and weigh about 250 mg of KClO3 into it (analytical balance). Melt the salt
and continue heating while O2 escapes in bubbles. At the end of the O2 evolution the melt solidifies
and can be molten again only with difficulty. Let the test tube cool for 2 – 3 minutes and complete
cooling under the water tap before weighing again. Calculate the molar mass of KCl from the
weight loss. Transfer a crystal of CrCl36H2O and 50 mg of KClO3 into another small test tube and
heat to melting. Observe the oxidation of Cr(III) to CrO42-.
Preparation of CuCl
The chloride of monovalent copper CuCl can be obtained by the reduction of Cu2+ wit elemental
Cu. If elemental Cu is added to a solution of CuCl2 the Cu becomes covered with a layer of
sparingly soluble CuCl which blocks further reaction. With sufficiently high chloride
concentration, however, attainable by addition of concentrated HCl or NaCl, CuCl is dissolved as
dichloro cuprate CuCl2- (see also the chapter about complex formation) and the redox reaction is
completed. Cu2+ exists as tetrachloro complex CuCl42- (greenish-yellow) in concentrated chloride
solutions. During the reaction the solution acquires deep brown hues (Cu(II) and Cu(I) together), at
62
the end the colourless CuCl2- is present. If the solution is diluted now, colourless CuCl is
precipitated.
4 g CuCl22H2O and 2 g elemental copper powder are weighed into a 50 ml conical flask and 4 g
NaCl are added. Add 20 ml concentrated hydrochloric acid, seal with a rubber stopper let the
reaction run on a magnetic stirrer. The decolouration of the solution marks the end of the reaction.
A large suction flask id filled with 150 ml deionised water and the small Buchner funnel equipped
with a filter paper is attached. Now the solution of CuCl2- which often is still contaminated with Cu
is sucked quickly through filter into the water in which colourless CuCl is precipitated. Change the
filter paper, pour the CuCl suspension into a wide-necked conical flask, clean the suction flask and
filter the CuCl. Wash once with water and three times with ethanol, suck to dryness finally.
Preparation of potassium tetrachloroiodate(III)
We produce a somewhat exotic state of iodine and illustrate the consequences of electronegativity.
The composition of potassium tetrachloroiodate(III) is, according to its name, K[ICl4]. Iodine has
oxidation state +III, chlorine, because of its higher electronegativity, -I and fulfils the octet rule.
The composite ion has a charge of -1, therefore the cation is a single K+. Starting materials are KI
and KClO3, which we know already. Additionally, the solution contains considerable amounts of
Cl- and H+, since we work in hydrochloric acid.
Reduction:
ClO3- + 6e- + 6H+
Cl- + 3H2O
Contrarily to the decomposition of KClO3 which was carried out in anhydrous melt we can no more
postulate O2- as a product, since this ion is more basic than OH-, the strongest base in water.
Therefore it is immediately converted to OH- by water, according to
O2- + H2O
2OH-
Since we even add acid there is immediate formation of H2O, as our reduction equation describes.
It is always important to consider the circumstances under which a process is carried out when we
formulate the reaction equations!
Oxidation:
I- + 4Cl-
ICl4- + 4e-
The required product is I+III bound to 4 chloride ions, we start at I-I, so 4 electrons are removed from
an iodine atom. In order to obtain the total reaction we have to find the smallest common multiple
of 4 and 6, again. This yields
63
2ClO3- + 12e- + 12H+
2Cl- + 6H2O
3I- + 12Cl-
3ICl4- + 12e-
combined to
2ClO3- + 12e- + 12H+ + 3I- + 12Cl-
2Cl- + 6H2O + 3ICl4- + 12e-
and simplified to
2ClO3- + 12H+ + 3I- + 10Cl-
6H2O + 3ICl4-
Ventilated hood!
2.03 g potassium iodide are weighed into a 100 ml conical flask and 10 ml of concentrated
hydrochloric acid are added. To this suspension a slightly warmed solution of 1.00 g potassium
chlorate in 5-6 ml water is added dropwise. A yellow solution is formed which yields golden
yellow needle-like crystals of potassium tetrachloroiodate upon cooling under the water tap. The
crystals are stable only in the solution not isolated therefore.
We have here, besides of some intermediate iodine, a redox process occurring in aqueous phase
which involves coordinative changes and completes only in strongly acidic medium. Write down
the reaction equation.
In qualitative and quantitative analysis redox reactions are used in many ways. Examples are the
permanganometric titration and the iodometric titration.
Redox reactions in qualitative analysis
a) The reaction of the silver ion Ag+ with thiosulphate S2O32- is used to detect S2O32-. Ag+ forms a
colourless sparingly soluble salt wit S2O32-, Ag2S2O3, which decomposes rapidly and the
coordinated sulphur of S2O32- remains as S2- under formation of black Ag2S with the Ag+ while the
rest of the thiosulphate coordinates the oxygen of a water molecule.
This example illustrates the purely formal character of oxidation states (they are indicated here in
Arabic numbers for clearness):
-2
O
-2 O
+6
O
O
S -2
S
-2
-2
-
-
-2 O
-
+4
S0
S
O
-
-2
64
The assignment to the two sulphur atoms is ambiguous. Based on electronegativity the oxygen
atoms are safely assigned -2. There should be no difference between two sulphur atoms which
supports the variant on the right. However, some people would prefer the left side variant because
of analogy to the sulphate ion. The total reaction
2Ag+ + S2O32- + H2O
Ag2S + SO42- + 2H+
is, based on the left side version, a pure acid/base reaction, however, based on the right side
formulation, also a redox reaction, since the central sulphur atom is oxidised from +IV to +VI while
the peripheral sulphur atom is reduced from 0 to –II.
Dissolve about 1 g Na2S2O35H2O in 20 ml water. One ml of this solution is added to 2 ml of
0.2 M AgNO3 solution and the reaction observed. In a second experiment add 4 ml thiosulphate to
1 ml AgNO3 solution. Observations?
With excess thiosulphate Ag+ forms a soluble complex Ag(S2O3)23-. This reaction is used for the
dissolution of unreacted AgBr from photographic films (fixation).
b) Dissolution of metal sulphides in nitric acid. There exist metal sulphides like NiS, CoS, which
cannot be dissolved in non-oxidising acids, e.g. hydrochloric acid. The reaction
MS + 2H+
M2+ + H2S
is too slow here. Dissolution can be effected, however, if the S-II of sulphides is oxidised by strong
agents like HNO3 to sulphate SO42-. The reduction of nitric acid yields NO and NO2 (red-brown
gas). Part of the sulphur can be precipitated as elemental S0 during the oxidation from S-II to S+VI.
Ventilated hood! Transfer some PbS or NiS into a large test tube, add some drops of concentrated
HNO3 and hest on the burner until the dark colour of the sulphide has vanished. If the material fails
to react add some more HNO3 and heat further. From PbS part of the Pb2+ will precipitate as
PbSO4(s) while NiS is dissolved completely. Prove the sulphate by addition of BaCl2 solution.
Write the redox reactions under the assumption that only NO is formed.
c) Separation of metallic mercury on copper metal. The standard reduction potentials for
Hg2+ + 2e-
Hg0 and Hg22+ + 2e-
2Hg0
are higher than the one for
65
Cu2+ + 2e-
Cu0
Therefore the reaction
Hg2+ + Cu0
Hg0 + Cu2+
runs to the right. Cu0 is said to displace Hg2+ from its solutions. If a blank copper wire is immersed
in a solution of Hg2+ (absence of strong complex forming agents presumed) Hg0 separates under
superficial alloy formation. By rubbing with a rag the silver-white mercury stain can be made
visible. After a while the Hg0 diffuses into the interior of the copper wire. The reaction is a
sensitive proof for Hg.
Dissolve a crystal of mercury(I) nitrate Hg2(NO3)22H2O in 1 ml water and acidify with a few
drops of 2 M HNO3. Dip a clean and blank splinter of copper into this solution.
d) Detection of chromium as chromate. Chromium(III) can be oxidised by hydrogen peroxide H2O2
to CrO42-. This can be recognised by the yellow colour or detected as yellow PbCrO4(s), BaCrO4(s)
or as brick red Ag2CrO4(s).
Dissolve a few crystals of the green CrCl36H2O in 2 – 3 ml water and make the solution alkaline
by adding 2 pellets of NaOH. Green Cr(OH)4- is formed. Add some drops of 10% hydrogen
peroxide and boil. After acidification of the solution with acetic acid the chromate can be
precipitated as yellow barium chromate BaCrO4 by addition of barium chloride solution (in the
automatically formed acetate buffer).
e) Redox system Fe(II) / Fe(III) - Sn(II) / Sn(IV)
Fe3+ can be reduced to Fe2+ by Sn(II) under formation of Sn(IV):
2 Fe3+ + SnII
Fe2+ + SnIV
In order to avoid hydrolytic effects we have to work in weak hydrochloric acid solution. The
completion of the reaction can be followed by the addition of some thiocyanate SCN-. This forms
deep red-brown complexes Fe(SCN)x(3-x)+ the colour of which disappears upon complete reduction
of Fe(III).
66
In a large test tube some crystals of iron(III) chloride FeCl36H2O are dissolved in dilute
hydrochloric acid and a few crystals of ammonium thiocyanate NH4SCN are added. Add a solution
of tin(II) chloride SnCl22H2O in dilute hydrochloric acid dropwise until the iron containing
solution is discoloured suddenly.
f) Proof of oxidising agents. Conversion of I- into I2
Oxidising agents can be detected by letting them act on iodide I- in acidic solution. The iodine
formed can be recognised by its brown colour. An example is iodate IO3-. If solutions of potassium
iodide and potassium iodate are mixed no colour results. I2 is not formed because the necessary
protons to complete the reaction (write an equation!) are missing. If acid, e.g. 2 M HCl is added the
reaction occurs immediately. With KIO3 and KI, or KBrO3 and KBr quantitatively known amounts
of I2 or Br2 can be prepared.
Carry out the reaction previously described with iodide.
Disproportionation of H2O2, catalase
Hydrogen peroxide H2O2 is a water-like substance, however strongly oxidising because it can be
reduced to water. On the other hand H2O2 can also be oxidised to O2. The oxidative power is so
great that this process, called disproportionation, really occurs at a slow rate and makes H2O2
unstable on the long term. Write the corresponding reaction equation. H2O2 is formed e.g. during
the biological reduction of dioxygen, one of the principal energy sources of life. Since H2O2 is toxic
due to its oxidative power nature has evolved biocatalysts (enzymes) which accelerate the
decomposition reaction substantially. These enzymes are built around copper(II) ions. Copper(II)
alone catalyses the reaction, like many transition metals. The best known enzyme in the group is
catalase. We show here the activity of Cu2+ and catalase against H2O2.
Prepare 10 ml of 0.1 M CuSO4 solution and add 1 drop of concentrated NH3; prepare also 10 ml of
1 M Na2S and peel a potato. 5 g of this a re cut into small pieces and mashed in a mortar. The mash
is mixed with 10 ml water in a small beaker and allowed to stand for 10 minutes under occasional
67
stirring with a glass rod. (do never use metal, e.g. a spatula, for this). The extract is decanted and
centrifuged.
Transfer 2 ml potato extract into each of two test tubes, 2 ml Cu2+ solution into each of two others.
To one of the tubes with extract and to one with Cu2+ a drop of Na2S solution is added. To each test
tube 3 ml freshly prepared 1 M H2O2 are added. Observe. What is happening?
Standard reduction potential Fe(CN)63- / Fe(CN)64Standard reduction potentials are not easily measured in general because electrodes tend to be
unresponsive, means kinetically inhibited. Empirically it was found that certain additives which do
not show up in the reaction equations help to settle a potential faster. For the potential settling of
the system hexacyanoferrate(III) – hexacyanoferrate(II) at a gold electrode the addition of a trace of
Ag+ is the necessary "catalyst". This kind of difficulties is the major reason why redox titrations are
usually not followed by potential measurement at electrodes.
In this experiment the electrode potential of three solutions which contain Fe(CN)63- and Fe(CN)64in various ratios shall be measured at a gold electrode. The counter electrode is a calomel or
silver/silver chloride reference electrode with [Cl-] = 3 M, E = 0.200 V. The catalyst is a trace of
Ag+ ions.
Prepare solutions of potassium hexacyanoferrate(II) K4[Fe(CN)6]3H2O and potassium
hexacyanoferrate(III) K3[Fe(CN)6], 50 ml and 0.1 M each. From these solutions mixtures are
prepared which contain FeIII and FeII in the ratios 10:1, 1:1 and 1:10. Add a drop of 0.2 M AgNO3
to each mixture.
For the potential measurement a millivolt meter with an Au electrode and reference electrode is set
up ready for use. Fill the solutions in the vessels set up and read the corresponding potentials.
These have to entered into the Nernst expression with the potential of the reference electrode
corrected for the offset against the normal hydrogen electrode (E° = 0 V).
Fe(CN)63- + e-
Fe(CN)64-
RT [Fe(CN)63-]
E + Eref = E°(Fe(CN)63-/Fe(CN)64-) + F ln
[Fe(CN)64-]
68
Enter the values of E, [Fe(CN)63-], [Fe(CN)64-] into the equation and calculate
E°(Fe(CN)36-/Fe(CN)64-).
Permanganometric titration
Permanganate ion MnO4- is one of the strongest oxidising agents stable in aqueous solution, which
is applied in acidic (formation of Mn2+) as well as in neutral and alkaline solution (formation of
MnO2). In the presence of certain ligands the reduction of MnO4- can lead also to other oxidation
states, e.g. Mn(III). The intense violet colour of MnO4- is usually sufficient for end point detection
in titrations. Disadvantageous is only the capability of MnO4- to oxidise chloride to chlorine. With
its properties permanganate offers the possibility to titrate almost everything that can be oxidised in
water, I-, As(III), Sb(III), H2O2, VO2+, HOOC-COOH, NO2-, HS- etc. The individual methods can
be looked up in the appropriate analytical books like "Skoog & West: Fundamentals of Analytical
Chemistry" or "Arthur Vogel: Quantitative Inorganic Analysis".
Permanganometric determination of oxalic acid, (COOH)2
The carbon of oxalic acid H2Ox (Ethanedioic acid) is oxidised to CO2 by permanganate. The
reaction runs at a useful rate only in warm and acidic solution. The Mn2+ formed acts also as a
catalyst for the oxidation (autocatalytic reaction). Determine the valence of the carbon in oxalic
acid (according to the inorganic rules) and write down the complete titration reaction.
Execution
For the determination of the water content of oxalic acid (possible: 0, 0.5, 2 H2O or nonstochiometric) a sample of 120 to 140 mg of the oxalic acid is weighed with the analytical balance
and dissolved in a wide-necked conical flask in about 150 ml of water, under addition of about 7 ml
of concentrated sulphuric acid. Heat the solution to 60 °C and titrate with 0.02 M KMnO4 until a
slight pink-violet hue remains. Calculate the water content of your oxalic acid sample.
69
Iodometric titration (of Cu2+ solution)
In contrast to permanganate iodine I2 is a comparably weak oxidising agent, which is soluble in
aqueous solution in the form of triiodide
I3- + 2e-
3I-
With such a triiodide solution strong reducing agents like S2O32-, As(III), H2S, SO2 etc. can be
titrated directly. The indicator is soluble starch. This forms a deep blue-violet inclusion compound
with the free iodine.
This nice end point indication can be used in different ways. Oxidising agents stronger than iodine
can oxidise added iodide to iodine, e.g.
2Fe3+ + 2I-
2 Fe2+ + I2
This way the redox equivalents are quantitatively transferred to iodine which can be titrated with
standardised thiosulphate solution and starch as indicator from blue to colourless solution. The
thiosulphate is converted to tetrathionate:
I3- + 2 S2O32-
3I- + S4O62-
A selected iodometric determination is the one of Cu2+. It reacts with I- under formation of the
sparingly soluble iodide of monovalent copper CuI and the equivalent amount of I2
2 Cu2+ + 4 I-
2CuI(s) + I2
which can be titrated now with thiosulphate.
Execution of the iodometric determination of Cu2+. An aliquot of the Cu2+ solution handed out by
the teaching assistant is acidified to pH = 3-4 with some drops of acetic acid and an excess of solid
KI is added. CuI and triiodide I3- are formed. I3- is titrated with calibrated 0.1 M sodium
thiosulphate under formation of tetrathionate S4O62- and iodide. Short before the end point (solution
slightly yellow) a few drops of starch solution are added (boil 100 mg starch in 10-20 ml of water)
and the titration is finished with the change from blue-violet to colourless.
Preparation of the thiosulphate solution: dissolve 0.01 moles of Na2S2O35H2O in a little water,
transfer into a 100 ml graduated flask and fill to the mark. Add about 100 mg of sodium hydrogen
carbonate NaHCO3 for stabilisation.
70
Coulometric titration, analysis of S2O32- and H+
In a coulometric titration the titration reagent is generated by an electrode in an electrochemical
process, therefore the name. The yield of the reagent related to the current must be 100%, and the
reaction at the counter electrode must not interfere. If the current i is kept constant during the
titration the time needed to reach the end point is the quantity to be measured.
Today it is easy to obtain a constant current i ( 1 mA) and to measure this by means of
semiconductors. Similarly simple is the measurement of time of about 100 s with a precision of 1 s.
Under these conditions there are 0.001 in 100 s converted => 0.1 As = 0.1 Coulomb, which
corresponds to the charge of only 10-6 moles of electrons. Coulometry is a method suitable for
small amounts therefore. In order to avoid side reactions electrode compartments must often be
separated by diaphragms, e.g. glass filter frits. The following to examples are selected such that a
phase boundary is unnecessary (both electrodes in the same compartment).
Apparatus
The current generating part is a four piece block of 1.5 V batteries with a current controlling diode
of type CR 100 – 8238 which can deliver about 1 mA up to a cell voltage of about 3 V. The exact
current i an individual diode can be measured with an ampere meter. The .electrode system consists
of two platinum wires of 1 mm diameter and about 12 mm length, or of a platinum wire and a silver
wire. The titration vessel is a 10 ml narrow cylindrical bottle (14.5) equipped with a small stirrer
bar.
Circuit:
+
-
71
Titration of hydrochloric acid with coulometrically generated OH-.
Into a 10 ml narrow cylindrical bottle (obtain it at the equipment counter) about 200 mg solid KCl
are transferred, 1 ml of a 10-3 M HCl (Preparation: 1 ml 0.1 M HCl  100 ml) are added with a
pipette. A drop of phenol red indicator and 6 ml water complete the mixture. The electrode
combination is inserted and the apparatus is fixed on a stand, above a magnetic stirrer. The silver
wire serves as the anode (attached to the diode outlet) and the platinum wire is the cathode (minus
outlet of the battery). Under simultaneous start of time measurement the current is turned on and
the time is measured until the first deviation from pure yellow to red occurs.
Sources of error:
With such small amounts to titrate large amounts of indicator, dissolved CO2, adsorption at glass
surfaces etc. can lead to appreciable end point errors. Work meticulously and cleanly. It is possible
that the indicator solution contains already some Ind- besides HInd. In this case add first a drop of
indicator and acidify with some drops of 10-3 M HCl, neutralise at the cathode and finally add 1 ml
10-3 M HCl and "titrate".
Calculations:
Write the cathode reaction and calculate from i and t and calculate from i and t and the
electrochemical equivalent (1 Farady = 96494 Coulomb/mole) the amount of titrated hydrochloric
acid.
Write the anode reaction and calculate the amount of the product formed.
Titration of thiosulphate with coulometrically generated triiodide I3Transfer ca. 200 mg KI into the narrow cylindrical flask and pipette 1 ml 10-3 M thiosulphate
solution to this. The indicator is 2 drops of boiled starch. The mixture is diluted to 6 ml with water
and one drop of 100% acetic acid is added. Electrolyse between two platinum wires to the first
brown-violet hue.
Write the anode reaction and calculate, like before, the amount of titrated thiosulphate.
72
VII Ligand exchange and coordination chemistry
Introductory experiments, preparative coordination chemistry
Equilibria in complex formation
Metal indicators, complexometric titration
The designations "complex" or "coordination compound" are based on the latin words complexus =
embracing and coordinare = to assign. Originally it meant the deposition of molecules into a
compound under formation of a so-called higher order compound, e.g. the addition of NH3 to
copper sulphate.
CuSO45H2O + 4NH3
Cu(NH3)4SO4H2O + 4H2O
Today the name stands for a huge variety of compounds which contain "complex" particles,
particles which are built around a central atom with a number of nearest neighbour atoms, called
ligand atoms, the bonds being not purely ionic.
One kind of order follows the kind of central atom. They can be metals and non-metals, and the
next principle of order is their stochiometric valence. The kind of nearest neighbours, their number,
called the coordination number and the geometric arrangement, called coordination geometry, are
further characteristics.
Examples:
Central atom Valence Ligand atom Z Geometry
Hg(NH3)42+ tetrammin mercury
Hg
+II
N
4
tetraed
CuCl42- tetrachloro cuprate
Cu
+II
Cl
4
quadrat
AlF63- hexafluoro aluminate
Al
+III
F
6
oktaed
Ni(CO)4 nickel tetracarbonyl
Ni
0
C
4
tetraed
SbF5 antimony pentafluoride
Sb
+V
F
5
trig. bipyr
Ag(S2O3)23- bis-(thiosulfato) silver
Ag
+I
S
2
linear
The ligands can be simple atomic anions like F-, Cl-, O2- etc. The ligand atoms, however, can also
be part of a larger ion or molecule: N as ligand atom in complexing agent NH3, S as ligand atom in
73
SCN-, O as ligand atom in complexing agent H2O etc. A complexing molecule may contain more
than one ligand atom and be able to saturate more than one coordination position. For example,
ethylene diamine H2N-CH2-CH2-NH2, an uncharged complexing agent, can attach both its N atoms
to Cu2+ such that a five-membered ring, a so-called chelate ring (chele = crab claw) is formed. The
substances are called bidentate or polydentate ligands, respectively. The anionic polyphosphates of
earlier detergents were able to bind Ca2+ in chelate form, another example of polydentate
complexing agent. The interplay of central atom with ligand atoms is governed by certain bilateral
preferences. The electron configuration of the central atom plays the crucial role, and the
stochiometric valence (see chemistry lecture) is also important. The preferences are expressed in
selectivity rules which also can explain the order concerning the exchange of ligands. A more
firmly attaching ligand will displace one with a weaker bond from its position, or the ligand with
the higher concentration will displace the one with minor concentration by the mass action.
Exchange can occur only stepwise, however.
The complexes of the transition metals (electron configuration dq, 0 < q <10) are coloured. Colour
and colour intensity depend on the kind and number of coordinated ligands as well as on the
coordination geometry. Therefore, ligand exchange cannot only be followed visually, e.g. in
qualitative analytical determinations, but it can be measured also quantitatively with a
spectrophotometer. Together with the acidity/basicity of the free ligands and their changes by the
coordination a tool for the quantitative elucidation of reaction mechanisms of complex formations
is obtained, yielding stochiometries and equilibrium positions.
The rates at which transition metal ions exchange ligands is also dependent on the d electron
configuration (see chemistry lecture). Especially the central atoms Cr(III) (d3), Co(III) (d6), Pt(II)
(d8) and Pt(IV) (d6) show inert behaviour, means they exchange their current ligands only slowly
against others. The inertness of a metal centre generally increases with a more positive valence, and
from the first to the second and third transition row.
For the ligand exchange in aqueous solution two terms are rather common: the substitution of a
water (solvent) molecule in an aqua complex by a different ligand is called "complex formation"
M(H2O)xn+ + yL-
MLy(H2O)x-y(n-y)+ + y H2O
while the substitution of a ligand L by another ligand is called "ligand substitution"
MLx + y B
MLx-yBy + y L
74
Complex formation is only possible if the new ligand is more strongly coordinated than water. It
must be noted that water is a fairly good ligand itself, which can be seen in the order of magnitude
of the hydration enthalpies
M+
400 kJ Mol-1
M2+
1800 kJ Mol-1
M3+
4000 kJ Mol-1
Aqua complexes themselves are obtained by dissolution of a salt of the corresponding metal ion,
with anions that hardly tend to complex formation, like perchlorate, nitrate and eventually sulphate.
Further one has to consider that the metal aqua ions with high charge show acidic character in
water (e.g. Al3+) and are stable only in acidified solution.
As diverse the interplay of all facts influencing complex formation is, as diverse are its use,
applications and consequences. Some freely picked examples are metal ions encapsulated in
organic molecules like: sandwich complexes, a kind of complexes that facilitates the passage of K+
through cell membranes. Chelating agents which form sparingly soluble precipitates with metal
ions, useful in gravimetry. Water softeners to bind calcium. Complexones for the titration of
various metal ions. The red blood colour as an inert iron complex for the transport of O2. Vitamin
B12, a cobalt complex. Complex formations as detection methods in qualitative analytical
chemistry. Catalysts in polymer synthesis. Chlorophyll etc.
Introductory experiments in coordination chemistry
In these experiments a number of characteristic complex formations and ligand substitutions are
gathered, many of which are applied in qualitative and quantitative analysis.
a) Chromate CrO42- - chlorochromate CrO3Cl- - chromyl chloride CrO2Cl2
These three particles differ only, with identical coordination number, identical coordination
geometry and identical valence of chromium, by the numbers of ligand atoms O-II and Cl-I. CrO2Cland CrO2Cl2 can be prepared by ligand exchange reactions, starting with chromate and Cl-, under
protonation of the leaving ligand O2-. The use of 25% hydrochloric acid leads to CrO3ClCrO42- + Cl-
CrO3Cl- + H2O
75
If concentrated sulphuric acid is used for the protonation chromyl chloride, a volatile red-brown
material, is obtained. It is used in the detection of chloride.
CrO42- + 2Cl- + 4H+
CrO2Cl2 + 2H2O
If CrO3Cl- or CrO2Cl2 are brought into water the reactions are reverted.
1 g potassium dichromate K2Cr2O7 are dissolved in a mixture of 1 ml water and 1.5 g concentrated
hydrochloric acid under gentle heating. After cooling yellow-red crystals of KCrO3Cl begin to
separate. Filter on a glass filter frit and dry on filter paper in air. It can be proven that Cr and Cl are
present in the ratio of 1:1 as follows: Dissolve some crystals of KCrO3Cl in a small test tube in
water and add solid sodium acetate to generate a buffer. KCrO3Cl decomposes to CrO42- and Cl-.
Both can be precipitated selectively and centrifuged as silver salts according to the solids chapter.
What should be the ratio of Ag+ consumption for Cl- and CrO42-? Do the experiment!
b) Complex formations of Cu(II) and Fe(III) with Cl-, Br-, CH3COO-, SCN-, F-, NH3. By the
combination of solutions containing the aqua ions Cu2+aq or Fe3+aq with solutions of the ligands
mentioned the complex formations can be followed nicely by the occurring colour effects.
Prepare two solutions of 1 g iron(III) ammonium sulphate NH4Fe(SO4)212H2O and copper
sulphate CuSO45H2O each in 2 – 3 ml water. Add 2-3 drops of these solutions to the following
solutions, prepared in small test tubes (level about 1 cm high):
Tube 1
HNO3 2 M
Tube 2
HCl conc.
Tube 3
KBr 100 mg/1 ml H2O
Tube4
CH3COONa3H2O
Tube 5
KSCN
Tube 6
NH3 conc.
Add sodium fluoride solution dropwise to the thiocyanate complexes of iron(III).
c) Ammine complexes. Ammonia as complexing agent
NH3 forms soluble ammine complexes with many transition (Cu2+, Ni2+, Co2+) and B (Zn2+, Ag+)
metal ions, which are stable in excess ammonia and prevent hydroxide precipitation. Many highly
charged A metal ions (Al3+, Ti4+) prefer OH- as a ligand over NH3: addition of NH3 causes metal
hydroxide precipitation.
76
With Ag+ the ammine complex formation is such that AgCl is dissolved already in dilute NH3
while the more sparingly soluble AgBr is soluble only in more concentrated NH3, and the even less
soluble AgI cannot be dissolved even with concentrated NH3.
Three examples are selected to show the action of ammonia as a complexing agent:
Ni2+ + x NH3
Ni(NH3)x2+
To a solution of nickel nitrate Ni(NO3)26H2O which contains the nickel aqua ion an excess of
concentrated NH3 is added. Soluble blue-violet nickel ammine complexes are formed. The value of
x depends on the concentration ratios: x can be 4, 5 or 6.
Cd(OH)2(s) + 4NH3
Cd(NH3)42+ + 2 OH-
Dissolve some crystals of cadmium sulphate 3CdSO48 H2O in 1 ml water and precipitate the
cadmium by addition of 2 M sodium hydroxide as gelatinous cadmium hydroxide Cd(OH)2. After
addition of concentrated ammonia the cadmium is re-dissolved as a tetrammine complex.
AgBr(s) + 2 NH3
Ag(NH3)2+ + Br-
Dissolve some crystals of KBr in 1 ml water and precipitate Br- with 0.2 M silver nitrate. The
yellowish AgBr precipitate can just be dissolved with concentrated NH3.
d) Formation of hydroxo complexes of Al(III) and Zn(II)
Most metal hydroxides are sparingly soluble. Some of these hydroxides, e.g. Al(OH)3 and Zn(OH)2
are dissolved by addition of OH- as hydroxo complexes (aluminate, zincate). This behaviour is
called amphoteric.
Al(OH)3(s) + OH-
Al(OH)4-
Zn(OH)2(s) + 2 OH-
Zn(OH)42-
Sodium hydroxide solution is added dropwise to solutions of aluminium chloride and zinc sulphate.
Observe the prcipitation of the hydroxides and the following re-dissolution under formation of the
hydroxo complexes. For Al(OH)3 0.1 M sodium hydroxide is sufficient for complex formation, for
Zn(OH)2 a higher OH- concentration is required, 2 M NaOH. Sodium hydroxide can even dissolve
aluminium metal directly to aluminate. Write the reaction equation.
77
e) Formation of iodo and thio complexes of Hg(II)
Hg(II) has a prominent preference fot I- and S2- as ligands. The red HgI2 itself is sparingly soluble
in water but is dissolved readily upon addition of I- as colourless iodo complex. The complex
formation is so extended that iti is possible to titrate HgI2 directly with KI solution. At the end point
the red precipitate disappears.
HgI2(s) + 2I-
HgI42-
Similarly HgS is sparingly soluble but dissolved as thio complex HgS22-, a property used in
qualitative analysis to separate Hg(II).
HgS(s) + S2-
HgS22-
Remember that e.g. CuCl forms a chloro complex CuCl2- with excess chloride (see redox chapter).
Tare a large test tube, weigh about 0.5 g KI exactly and dissolve in about 3 ml water (note the
weights: tube, KI, solution). Into a second tared test tube weigh about 0.5 g HgI2 exactly. Dissolve
this by dropwise addition of the KI solution. Shake well after each addition. The end of the reaction
is recognised by the total disappearance of the red solid. Weigh the test tube and calculate the
coordination number of the iodo complex from all data collected.
For the formation of the thio complex HgS22- transfer a few crystals of HgCl2 into a test tube and
dissolve in 2-3 ml water. Add a solution of Na2S9H2O dropwise. Initially the sparingly soluble
HgS is formed which dissolves again by the addition of more sulphide as HgS22- (effectively HS- is
added).
Inertness of Fe(CN)64The complex particle Fe(CN)64- should decay into Fe2+ and hydrocyanic acid upon acidification:
Fe(CN)64- + 6 H+
//
Fe2+ + 6 HCN
This decomposition is very slow, however, Fe(CN)64- is – like Fe(CN)63- - an inert complex.
Therefore one can isolate the corresponding free acid upon acidification.
Fe(CN)64- + 4 H+
H4[Fe(CN)6]
78
The acid forms a sparingly water soluble addition compound with ether which allows for the
isolation. The dry acid is stable for unlimited time while in moist air it is slowly decomposed under
blue colouration (Prussian blue).
2 g K4Fe(CN)6 are dissolved in 18 ml water and 5 ml concentrated hydrochloric acid are added.
The KCl initially precipitated is just dissolved by adding water dropwise. Now 2-3 Pasteur pipettes
of ether are added which causes the ether adduct of H4[Fe(CN)6] to precipitate as leaf like crystals.
These can be filtered on a glass filter frit. Purification is possible by dissolution in ethanol and recrystallisation with ether. This purification shall not be carried out here.
g) Stepwise complex formation, iron complexes of tiron
Tiron is an ortho-diphenol with two -SO3- functions incorporated to improve solubility
in water. The two phenolic -OH functions are responsible for its property to act as a bidentate
chelating agent. A metal with the coordination number KZ = 6, e.g. Fe3+,
-O 3S
OH
-O 3S
OH
can bind totally three of these tiron anions L2-. In the formation process, according to
Fe3+ + H2L
FeL- + 2H+ etc.
two protons are released per L. Thus the pH of the solution has an influence on the complex
formation. With Fe(III) the conditions are such that, excess of ligand presumed, in acidic solution
(pH = 2) only the blue 1:1 complex FeL is present, at pH  7 the violet 1:2 complex FeL25dominates and at pH  10 the red 1:3 complex FeL39- prevails.
Into a large test tube the smallest possible trace of FeCl36H2O is transferred, dissolved in 10 ml
water and acidified with one drop of 2 M HCl. Add a spatula tip of tiron. The blue 1:1 complex
appears. By addition of solid CH3COONa3H2O the pH is increased until the colour changes to
violet (1:2 complex). Upon addition of a drop of concentrated NH3 the red 1:3 complex is formed.
It can be shown that the bidentate complexing agent tiron is bound more strongly than a
monodentate one since addition of F- causes no discolouration as it was the case with the
thiocyanate complexes of Fe(III).
79
Preparative coordination chemistry
Many complex particles are also known to occur in solid compounds. Such solids are not only
proof for the existence of the complex with the corresponding stochiometric composition observed
in solution, but allow to obtain all structural details by means of X-ray diffraction analysis: precise
coordination geometry, bond distances, bond angles etc. Those solids are often easily isolated, the
preparation of previously unknown complexes demands, however, the profound knowledge of
chemical reactivity and perfect skills in laboratory methods. To isolate new compounds is one of
the fundamental goals in chemistry.
Two categories of solid coordination compounds are paid special attention here. The robust
complexes of Co(III), Cr(III), Pt(II), Pt(IV) etc. which exchange ligands only slowly have enabled
the preparation of the various isomers of their coordination compounds. For example, cisCo(NH3)4Cl2+ and trans-Co(NH3)4Cl2+ can be isolated separately. Such solid isomers are the ideal
starting materials to study the kinetics of slow conversions.
Another group of solid coordination compounds is found in analytical applications. A central atom
can coordinate with single or double negatively charged bidentate ligands such that an electrically
neutral complex is formed. If this compound is stochiometrically uniform and sparingly soluble in
water it can be used for the gravimetric determination of the corresponding central atom. Typical
examples are 8-hydroxyquinoline, cupferron, dimethyl glyoxime etc.
On the following pages there are two methods for the preparation of typical representatives of solid
coordination compounds. The list could be arbitrarily extended. Synthesise one of the two
compounds mentioned.
Tetraammine nickel nitrite Ni(NH3)4(NO2)2
This is a nickel complex with only 4 NH3 bound to nickel. The two nitrite ions which compensate
for the charge of the central ion are doubtlessly bound to the two remaining coordination positions
of Ni. Noticeable is the deep red colour of this complex. Compare with the description of Ni(CN)42.
80
O
-
O
N
H3N
NH3
Ni
H3N
++
NH3
N
O
O
-
6.2 g nickel acetate Ni(CH3COO)24H2O are dissolved in as little water as possible under gentle
heating. In order to avoid hydrolysis a drop of concentrated CH3COOH is added. 30 g
CH3COONH4 together with 20 g sodium nitrite are dissolved in a 300 ml conical flask, also with as
little water as possible. Now add the nickel acetate solution and further 15 ml concentrated NH3.
After a while the red complex is precipitated in the form of delicate crystals, sometimes one has to
wait overnight. The supernatant solution is decanted and the crystals are transferred into a Buchner
funnel with filter paper by means of ethanol as the washing fluid. Dry at room temperature on the
filter paper.
Potassium dioxalato cuprate(II) K2Cu(OOCCOO)22H2O
The oxalate ion is a complexing agent which forms sparingly soluble compounds like FeC2O4,
La2(C2O4)3 etc.
O
O
-O
O
-
Oxalate can be coordinated up to quadridentate, e.g. in calcium oxalate CaC2O4(s). If such oxalates
are treated with more oxalate often anionic oxalate complexes are formed in which the oxalate is a
bidentate ligand like in La(C2O4)33-, Co(C2O4)33-. This kind of complex can often be isolated as an
alkali salt.
Prepare a hot solution of 7.3 g dipotassium oxalate with 20 ml water in a 100 ml conical flask. In a
100 ml beaker 2.5 g copper sulphate CuSO45H2O are dissolved in 10 ml water under heating. Add
a few drops of the hot potassium oxalate solution. Bright blue CuC2O4(s) is precipitated
immediately. Add the rest of the oxalate solution until the copper salt is completely re-dissolved
(magnetic stirrer). Upon cooling blue K2Cu(C2O4)22H2O crystallises, it is filtered on a glass filter
81
frit of degree 4. Wash with about 5 ml water, then with about 10 ml ethanol and finally with 10 ml
ether. Dry on a filter paper in air.
Metal indicators
Metal indicators are a combination of dye and bi- or polydentate ligand. When they bind to metal
ions their colour changes, exactly like a pH indicator that shows the uptake or loss of H+ by a
colour change. In fact metal indicators are also pH indicators. The main application of metal
indicators is in analogy the titration of metal ions with the solution of a chelating agent (mainly
EDTA), where a large change of pM = -lg[M] takes place at the end point and the metal indicator is
forced to change its colour with the pM change. Not every metal indicator is suitable for every
metal. The metal indicator must not bind more strongly to the metal than the titration reagent or the
end point colour change
M(Ind) + Y
MY + Ind
does not take place, at least not at the expected concentration. The pH indicator property must be
eliminated, by working in buffered solution at constant pH.
A further application stems from the often high extinction coefficients of many M(Ind) complexes:
the concentration of M(Ind) can be measured optically and is therefore useful in analysis. With
them it is possible to determine metal ion concentrations as small as 10-8 M. Dithizone is such a
complexing agent which helps to detect tiniest amounts of "heavy metal" ions.
Some characteristic metal indicators are listed below, without drawings of the complicated
structures, together with their indicating properties. Further details can be found in specific
brochures and analytical textbooks (Vogel: Quantitative Inorganic Analytical Chemistry).
Study the colour changes of the following indicators. A highly dilute solution of each indicator is
adjusted to the pH indicated. Add a small quantity of the correspondingly listed metal salt to each
of the indicators.
Dye
pH/(reagent)
Metal
Murexide
14 (NaOH)
Cd2+
Erio T
9-10 (NH3/NH4+)
Ca2+
Pyridyl-azo-naphthol (PAN) 5 (CH3COOH/CH3COO-) Cu2+
82
Note the colours before and after metal addition in a table. Finally, add 0.1 M Komplexon III
solution dropwise to each of the solutions that had metal added. What happens, and why?
Dithizone as metal indicator
Transfer about 5 ml dichloromethane CH2Cl2 into a large test tube and dissolve a trace of dithizone
in it. The solution shall be green. Pour a layer of 20 ml deionised water on the solution in which a
single small crystal of ZnSO47H2O has been dissolved previously, and shake. Zn2+ forms a
dithizone complex which is soluble in CH2Cl2 with violet colour.
Complexometric titration
For a polydentate ligand which saturates multiple or even all of the coordination sites of a metal ion
a very simple stochiometry of the complex compound is expected, namely 1:1. The only
complications are the coordination of H+ to the ligand or the coordination of OH- to the metal ion.
The occurrence of a precise 1:1 stochiometry
ML + iH+
M + Hi L
can be explored for analysis by the titration of a metal ion with such a chelating agent if one
succeeds in the detection of the end point with a metal indicator. In practical use we have today
only the anion of ethylene diamine tetraacetic acid
O
O
CH2
HO
+
NH
O
C
O
-
OH
OH
C
CH2
CH2
CH2
HO
C
O
+
CH2HN
O
CH2 C
+
NH
O
O
+
HN
-
O
O
-
-
O
O
H4EDTA, abbreviated H4Y
83
as a reagent of significance. It is a hexadentate ligand if fully coordinated. The disodium salt
Na2H2Y2H2O is commercially available under a variety of names (Komplexon III, Sequestren,
etc.).
Preparation of the calibrated Komplexon III solution, 0.1 M
Weigh exactly 37.225 g Komplexon III, dissolve in deionised water and fill in a graduated flask to
1000 ml. For the most precise work the solution must be calibrated since the crystal water content
of Komplexon III is not entirely reliable.
Titrate 10 ml quantities of a solution of about 1 g exactly weighed CaCO3, which is prepared in a
100 ml graduated flask, with the Komplexon III. The preparation of the Ca2+ solution is somewhat
tricky. The CaCO3 is not water-soluble, but can be weighed precisely. It is transferred with little
water into the flask and dissolved by dropwise addition of 6 M HCl and shaking until the solid is
just dissolved. Continue shaking until the CO2 evolution ceases. Fill to the mark with water and
mix. For the titration add 5 ml of the ammonia buffer described below to the 10 ml quantity, and a
spatula tip of Erio T ground with NaCl 1:100.
Why is the buffer required?
Determination of the hardness of water by complexometric titration
50 ml tap water (2 x 25 ml pipette) or 5 ml mineral water (do not use oligomineralised mineral
water! This designation is a bad joke anyway! Examples: Evian, Arkina, Henniez, Vichy, Badoit,
Perrier and other life style waters. Suitable are Aproz, Passuger, Rhäzünser, San Pellegrino, Valser,
Lostorfer, Meltinger etc.) are transferred into a wide-necked conical flask (add 40-45 ml deionised
water to the mineral water). To the solution 2-3 Pasteur pipettes of a buffer which is prepared from
7 g NH4Cl, 57 ml concentrated NH3 and 43 ml H2O are added. The pH should be 9-10 (check with
glass rod and indicator paper). Now a spatula tip of ground Erio T/NaCl 1:100 is added and the
titration with 0.01 M Komplexon III (= EDTA2 H2O = Na2H2, dilute the previously prepared
solution by a factor of 10) is carried out until the colour changes from red to pure blue, use a white
paper as background. Check the pH from time to time and add buffer if the pH becomes lower than
9. Do two titrations, better three, and average. If the titration volume is above 50 ml (mineral
84
waters or tap water from limestone areas) one should consider the use of higher EDTA
concentrations (0.05-0.1 M).
Evaluation: calculate first [Ca2+ in Mol/l in the sample, then use the following conversion factors:
French degrees (frequent on detergent packs)
[Ca2+]  MW (CaCO3)  100 = fH
(1 fH = 10 mg CaCO3 per litre)
German degrees
[Ca2+]  MW (CaO)  100 = dH
(1 dH = 10 mg CaO per litre)
very soft
0-4 dH
soft
4-8 dH
medium hard 8-12 dH
rather hard
12-18 dH
hard
18-30 dH
very hard
> 30 dH
85
IX Chromatography and liquid/liquid distribution
Chromatography
Chromatographic methods utilise the differences in adsorption of polar molecules on polar surfaces
of solids for separations. The equipment consists of a stationary solid phase on which the particles a
are deposited for limited intervals, and of a mobile phase which carries them along from deposition
to deposition.
The different adsorptive properties of different particles cause their different average deposition
times and therefore the separation. If liquids are used as mobile phase the adsorption is also
determined by the polarity of the liquid, and this is in turn determined by its composition (e.g.
mixtures of methanol and acetone). The forms of realisation are many: the stationary phase can be
stacked in a tube to form a column, called column chromatography; a piece of filter paper as
stationary phase: paper chromatography; a glass or plastic plate covered with a layer of a powder:
thin layer chromatography, etc. Even gases can be used as a mobile phase; the stationary phase in
this case is a viscous non-volatile organic fluid which is distributed in a thin layer over the inner
surface of a quartz or metal tube.
Chromatography is mainly a field organic chemistry (applications: isolation, purification,
preparative separation). Our first experiment demonstrates the separation of dyes with thin layer
chromatography.
Chromatographic separation of dyes
(Thin layer chromatography, TLC, on silica gel plates)
On the one of the narrow sides of a rectangular commercial TLC silica gel plate colour dots are
placed with 5 different felt tip pens about 5 mm from the edge, spaced by about 7 mm. For the
separation of the mixtures of colours (development of the chromatogram) a mixture of 1 part of
ethyl acetate, 1 part of methanol and 2 parts of acetone is transferred into a 400 ml beaker such that
the bottom is covered 2 mm high. The TLC plate is placed upright into the solvent, dots at the
bottom side. Cover the beaker with a filter paper or a Petri dish. The solvent is slowly soaked up by
86
the silica gel and the individual dyes are transported upwards at different velocities. When the
solvent front has reached the upper edge of the plate it is removed from the beaker and allowed to
dry in air. The relative run length Rf of a dye ist its absolute run length divided by the run length of
the solvent front. Under fixed physical conditions (temperature, partial pressures, composition of
solvent) this value is characteristic for the dye and allows for comparison of the compositions of the
colours.
TLC plate
Eluent front
Run lengths of
components
Start line
Liquid-liquid distribution
This separation method relies on the distribution of a substance A between two solvents in contact,
however immiscible. For example, if the solvents water and carbon tetrachloride CCl4 are used, and
as substance A the salt KBr is added, the salt has a much greater affinity to water than to CCl4. The
KBr added will be found almost exclusively in the water phase. On the other hand, if a material of
non-polar molecules is added as substance A (e.g. elemental bromine Br2) to the solvent system, it
will be found mainly in the non-aqueous phase, CCl4. This offers a possibility to separate salts from
molecular substances. In combination with suitable chemical reactions these preferences for one
solvent lead to very elegant separation methods. Charged particles which contain a group or atom
of type B, for example, can be transformed into uncharged particles containing B by a chemical
87
reaction. These can then be separated from other ions by extraction into an organic phase.
Examples for such conversions from charged to uncharged particles are
2 BrCH3COO- + H+
Al3+ + 3Hoxin
Br2 + 2eCH3COOH
Al(Oxin)3 + 3 H+
oxidation
protonation
complex formation
Salts which consist of very large ions are often soluble in organic solvents and can be extracted
from an aqueous phase into an organic phase. Co2+ e.g. can be extracted from water into ether in
form of the voluminous (NH4)2Co(NCS)4 (ion pair formation).
The solvent extraction method has found two major applications. On one side small amounts of
substances in large volumes of aqueous solution can be easily transferred into a small volume of
organic solvent and be concentrated thus. This is used for the extraction of organic acids from
aqueous solution into ether. Small amounts of radioactive materials can be removed easily and
quickly from aqueous solutions. Another application allows for the separation of mixtures of metal
ions if the individual metal ions show a distinguished difference in complex stability with a certain
ligand, e.g. Cl-. A metal that forms an uncharged chloro complex more easily will be transferred
preferably into the organic phase. Historically this way it was possible to separate the elements
scandium and thorium which can be hardly isolated from each other by other methods.
In the next experiment the distribution coefficient of elemental iodine kI for the system
water/dichloromethane CH2Cl2 is determined. In both phases iodine exists as I2. For the
determination of kI am small amount of I2 is equilibrated between H2O and CH2Cl2 before the I2
concentration is analytically measured in both phases.
Determination of the distribution coefficient of iodine in the solvent system H2O/CH2Cl2
Attention: CH2Cl2 is not totally harmless! Always use a rubber balloon to pipette CH2Cl2 and avoid
skin contact with CH2Cl2 and the inhalation of CH2Cl2 vapours. Used mixtures containing CH2Cl2
are to be poured into the dedicated waste containers.
88
10 ml of a 0.02 M I2 solution in CH2Cl2 are pipetted into a 300 ml conical flask filled with 200 ml
deionised water. A magnetic stirrer bar is added and the flask is sealed with a piece of aluminium
sheet. Stir for 15 minutes. After separation of the phases 150 ml of the aqueous phase are sampled
with a graduated cylinder and transferred into a 250 ml wide-necked conical flask. The iodine is
titrated with 0.01 M thiosulphate solution, according to the method in the redox chapter.
Under the conditions of our experiment the distribution coefficient is calculated according to
[I2]H2O
kI2 = [I ]
2 CH2Cl2
with
[I2]H2O =
n(I2)H2O = 0.2 l  [I2]H2O
n(I2)tot - n(I2)H2O
0.01 l
0.01MVtitr
2150 ml
[I2]CH2Cl2 =
and
n(I2)tot = 0.05 M  0.01 l
Ion exchangers
Ion exchangers make available a kind of chromatography for the separation of ions. They are solids
or polymers with cavities that can be filled with water. On the inside of the cavities there are
charged atoms groups the charges of which must be neutralised by counter ions in the aqueous
phase. These ions in the aqueous phase are mobile and can be exchanged for different ions of the
same charge type.
The most important types of ion exchangers are:
a) Inorganic aluminium silicates (zeolithes etc.), having a silicate lattice bearing negative charges.
Serve as cation exchangers.
b) Organic polymers with covalently bound –SO3- groups: R-SO3-Na+ are cation exchangers; with
covalently bound –NR3+ groups: R-NR3+Cl- are anion exchangers.
Ions attach with different strengths at the ion exchanging resin, generally the better the higher the
charge and the smaller the radius of the hydrated ion is. Thus the following series result:
89
worse
Li+< H+ < Na+ < Cs+ < Mg2+< Ca2+< Al3+< Ce3+
F-< Cl-< Br-< NO3- < HSO4- < I-
worse
better
better
If an ion exchanging resin RA+ were transferred into a solution of ion B+ and thoroughly mixed, the
equilibrium
RA+ +
B+
RB+ + A+
would be established at one single level. If the resin is, however, piled up in a glass tube to form a
column and the solution of B+ is passed slowly through it the equilibrium is established multiple
times at changing levels over the whole length of the column which results in a great separation
effect. Since the column starts with pure RA+ separations are possible even with unfavourable
equilibria because of the mass action.
An example is the preparation of a solution of rhodanic acid HNCS which is not stable in free form.
This is done by passing a solution of KNCS over a cation exchanger in acidic form RH+ and swaps
the potassium ion K+ for H+. The following experiment is an example for the separation of
transition metal ions of the 3rd period in different oxidation states and in the form of different
chloro complexes on an anion exchanger. The separation does not succeed with commercially
available cation exchangers.
90
Ion exchange chromatography: separation of Cu2+, Ni2+, Fe3+
Chromatography
tube
Mobile phase
Separating agent
(here: ion exchanging resin)
Glass wool
The separation of these three cations relies on the different properties of their chloro complexes.
Ni2+ has a poor affinity to Cl- and forms only NiCl+ in 9 M HCl, while Cu2+ and Fe3+ forms tri- and
tetrachloro complexes. The charge density which determines the bond strength with the ion
exchanger is different for CuCl3- and FeCl4- because the smaller copper complex has a larger radius
in the hydrated form and binds more weakly.
Method: the chromatography tube is mounted vertically on a stand and a ball of glass wool is
inserted just above the stopcock. 20 ml water and 5 ml 8 M HCl are added to 20 ml of anion
exchanger (Amberlite IR-400 or a similar one of the Amberlite IR-4xx series) in a 100 ml beaker
and stirred. The mixture is poured at once into the tube (stopcock slightly open, place beaker under
outlet). Rinse the beaker with 3 ml 8 M HCl and pour also into the tube. Knock gently at the tube,
91
best with the wooden clamp, such that the resin settles and air escapes. Never let the column run
dry, or it has to be refilled since air cannot easily be removed. Insert another ball of glass wool at
the top of the column. Flush the column with 15 ml 8 M HCl to ensure that it is in the Cl- form.
Close the stopcock when the HCl level is 1 cm above the glass wool. Add 0.5 ml of cation mixture
handed out by the teaching assistant to the column top. Add 1 ml of 8 M HCl and drain the column
until the upper level is again 1 cm above the glass wool ball. Add 15 ml of 8 M HCl to the column
top and open the stopcock cautiously until the outflow is 1 drop per second. Collect the outflow in
4 ml portions in test tubes (enumerate!). When the top level has reached again 1 cm, add 30 ml 3 M
HCl and continue collecting. As soon as this solution is also drained, add 30 ml water and collect
until you have a total of 20 test tubes with 4 ml eluted solution each. Regenerate the resin by
successive treatment with 10 ml 1 M Na2SO4 and 30 ml H2O, remove it from the tube and return it
to the teaching assistant.
Analysis of the samples: take 6-10 drops from each test tube for each test and examine as follows:
Cu2+: make slightly basic (pH=9) with concentrated NH3. Add some drops of 2,2'-bisquinoline or
neocuproin (2,9-dimethyl-1,10-phenanthrolin) solution (in 96% ethanol) and 4 drops of
hydroxylamine hydrochloride or hydroxylamine sulphate solution. Cover with some ether and
shake. If the ether is coloured blue-violet or reddish respectively, Cu2+ is present.
Fe3+: add 5 drops of 1 M KSCN. A distinct red colour indicates Fe3+.
Ni2+: make the solution alkaline by adding conc. NH3. Add some drops of 1% dimethyl glyoxime
solution (in 96% ethanol). A red precipitate formed within 10 minutes indicates Ni2+.
92
X Qualitative analysis
Qualitative analysis – an exercise
The knowledge about chemical reaction types acquired in our lab course until today can be used,
together with characteristic substance properties, to separate mixtures into components and to
detect their presence. In many cases it is impossible to detect a component in a mixture directly and
selectively, usually a preceding separation is required. Only the progress of the separation
together with a detection reaction at the end can prove the presence of a certain substance,
the direct application of this reaction often yields unclear results, since several components react
similarly.
We shall examine a mixture of 6 inorganic ions as an example, this fits the other subjects in the
course. The mixture contains
Fe2+, Zn2+, Ca2+, Cl, Br, I
These are ions which occur in pharmaceutical preparations, though usually not all together, but it
adds a reality touch to the exercise. It is typical for analytical chemistry that only certain
components in a sample are searched for. The idea of a total analysis of an environmental sample or
a natural material is completely unrealistic, since they might contain thousands of components.
Even with simpler problems the investigation is normally limited for economic reasons. In this
respect our approach is less realistic.
Before one starts an analysis, one has to know the characteristic chemical properties of the
components under examination. According to them the separation method is selected.
Fe2+: weakly colored transition metal ion, stable in acidic solution, becomes oxidized by air to
yellow-brownish Fe3+ in neutral to alkaline solution. It precipitates with hydroxide in alkaline
solution. Fe3+ forms many colored complexes
Zn2+: colorless (d10 !) transition metal ion, soluble as hydrated cation in acid and as hydroxo
complex in alkali, it is precipitated as Zn(OH)2 at neutral pH. This behaviour is called amphoteric
93
and is also observed with other metals, e.g. Al3+ and Sn2+. No remarkable redox chemistry except
for the reduction to Zn0. Zn2+ forms complexes which can be colored, but only by the ligand alone.
Ca2+: colorless earth alkali metal ion, shows a red color in the burner flame. A red flame color can
also be caused by Li+ and Sr2+, therefore the observation is only a hint, but not a proof. The ion is
water-soluble at most pH values except for strongly basic conditions, where is precipitated as
Ca(OH)2, and also in the presence of most monovalent anions except for fluoride. Anions with
higher charge like sulphate, carbonate and phosphate all form sparingly soluble salts with calcium.
Cl, Br, I are, besides fluoride, the halides, and the have very similar chemical behavior. All ions
are colorless, since the have a filled valence shell (noble gas configuration). All of them form
sparingly soluble silver compounds, which are colorless (Cl), light yellow (Br) or yellow (I). If a
precipitate with Ag+ has a pure white appearance, there can be only Cl, if even slight yellow
shades are visible there are many possibilities. The largest difference in the halides is there redox
chemistry: iodide can easily be oxidized to the elemental state, I2, with bromide it also possible, and
with chloride it is difficult. Since iodine and bromine, the elements, are colored, the oxidation can
be used for their detection.
Before one starts with separations and detections, a set of reagents and equipment should be
prepared in order to allow for fluent work. Our small problem requires only a reduced set compared
to a classic separation sequence.
Preparation of the artificial sample: Similar amounts of FeCl2, CaCl2•2H2O are mixed with half the
amount of ZnBr2 and ZnI2. All chemicals required are available at the department shop.
Homogenize thoroughly with mortar and pestle. 200-300 mg of the mixture are needed per student.
Equipment: spatula, glass rod, large and medium test tubes, small beakers of 50 and 100 ml,
centrifuge tubes, pH indicator paper (better continuous roll than sticks), magnesia rods, Pasteur
pipettes.
Apparatus: centrifuge, water bath (250 or 600 ml beaker containing de-ionized water, boil on
electric heater, then set thermostat to 120-140 °C, cover with a watch glass when not in use to
reduce vapor loss), gas burner.
Chemicals: solutions ready for use are: conc. sulphuric acid, conc. acetic acid, hydrogen peroxide
30%, hexane or light petroleum.
94
Solutions to prepare: hydrochloric acid 6 M (half conc.), sodium carbonate 2 M, sodium hydroxide
2 M, silver nitrate 0.1 M, K4[Fe(CN)6 0.1 M, Co(NO3)2 1% (weight).
Solids: potassium or sodium nitrite, potassium permanganate, ammonium chloride, potassium or
ammonium thiocyanate.
Procedure: all observations have to be listed in the laboratory journal. Try to formulate reaction
equations! Identify the according reaction type. We start with
-
Flame color: transfer a small amount of the mixture onto a watch glass by means of the
spatula tip. A magnesia rod is calcined in the non-luminous burner flame until the yellow
color ceases to fade. The cooled rod is moistened with 6 M HCl and dipped into the mixture
on the watch glass. Some material should be sticking to the rod now. The rod is carefully
introduced into the flame, preferably without dropping material into the burner tube.
Observe the flame color.
-
1. Separation: cations and anions are separated. Cations with two or more valences form
sparingly soluble carbonates which can be easily re-dissolved in dilute acid. Thus, all
cations except alkali ions can be separated elegantly from the anions. However, a
homogeneous solution of the sample has to be prepared first. This can be very difficult in
many cases. In our example it is quite simple: add about 3 ml H2O to 50-100 mg mixture in
a large test tube, which should dissolve most of the material after some shaking. Some
greenish flakes may remain. Complete dissolution is effected by dropwise addition of acetic
acid, whereas the solution is shaken well after each addition. Acid addition is stopped
immediately upon complete dissolution. This kind of reagent addition should always be
applied in order to avoid reagent excesses, which can be very confusing later and
increase the consumption of further reagents to be added. It is important that we do
not just execute a cooking recipe but act consciously to reach a goal. Afterwards,
sodium carbonate solution is added dropwise with vigorous shaking after each addition. At
the beginning, the solution might foam violently because of CO2 formation, caution!
Precipitation of carbonates sets in under strong clouding. As soon as we get the impression
that the precipitation is terminating we test for completeness: either we let patiently settle
the precipitate and add a drop of sodium carbonate gently at the solution surface, or we tilt
the test tube until we can see the tube wall through the liquid at the upper level, and add
sodium carbonate there; this is less precise. If, with the first or second method, clouding
appears at the contact position, precipitation is still incomplete. After completion the test
95
tube is stored for about 15 min. in the hot water bath in order to obtain more coarse grains
by re-crystallization. Finally, precipitate and solution are separated by centrifugation. The
carbonate precipitate is washed after decantation of the solution in order to remove
remaining anions in the slurry. This is done as follows: after decantation of the solution,
which is kept for anions determination later, 5 ml of H2O are added to the precipitate which
is the stirred up with the glass rod. Afterwards, the slurry is centrifuged and the supernatant
is discarded. The procedure is repeated once more. Now the precipitate is ready for cations
detection.
-
2. Separation: separate Fe2+/3+ from Zn2+ and Ca2+. The carbonate precipitate is transferred
into a beaker by means of a small amount of water and dissolved by dropwise addition of
acetic acid. Attention: strong foaming! Transfer the solution into a large test tube and expel
CO2 by gentle heating with the burner. Now adjust the pH to 10 with conc. NH3. Observe
meticulously and describe everything that happens here. At the end a precipitate of browngreenish flakes of iron hydroxides should be obtained. Ca2+ is not precipitated, the
hydroxide concentration is not sufficient. Zn2+ forms very soluble complexes with NH3. The
tube is heated for 10 min. on the water bath to improve the consistency of the precipitate,
which is finally separated by centrifugation. The solution is kept and the precipitate is
washed with a mixture of 0.5 ml conc. NH3 and 1 ml water this time, and the washing
solution is not discarded, but united with the solution obtained previously. This step is
repeated once more. The precipitate is kept for iron detection.
-
3. Separation: separate Zn2+ and Ca2+. The solution remaining from the previous separation
is boiled in a ventilated hood until it does no more smell of NH3. This can be done on a
stand in a beaker or in a large test tube. It is possible that a white precipitate appears. Now
the pH is brought to 4-5 with acetic acid, under dissolution of eventual precipitates. The pH
is made slightly alkaline again, to 8-9, by cautious addition of NaOH solution. A
voluminous mass of Zn(OH)2 precipitates. The product is again completed by heating on the
water bath for about 10 min. The precipitate is centrifuged and washed two times with 2 ml
water. The washing solutions are combined with the solution obtained after the first
centrifugation which should contain only Ca2+ now.
-
Detection of Fe as Fe3+: Fe2+ is not stable in air; therefore we conduct the detection for Fe3+.
The iron hydroxide precipitate is treated with 3 drops of 30% H2O2 and shaken thoroughly.
It should become distinctly brown now. The reaction is completed by heating for 10 min. on
the water bath. The precipitate is dissolved by dropwise addition of hydrochloric acid. As
soon as dissolution is complete we take a 0.5 ml sample of the solution and add
96
K4[Fe(CN)6] solution, which results in a dark precipitate of Prussian Blue (see chapter II).
Another 0.5 ml are mixed with a few crystals of solid KSCN or NH4SCN, which produces a
deep red complex (see chapter VIII).
-
Detection of Zn2+: the precipitation at neutral pH is already a strong hint for the presence of
this amphoteric element. We dissolve half of this precipitate by dropwise addition of HCl
and add K4[Fe(CN)6] solution which should produce an almost colorless precipitate:
K2Zn3[Fe(CN)6]2. This can be re-dissolved only with conc. HCl. Another, technically more
difficult method: the other half of the Zn(OH)2 precipitate is placed on a watch glass. A
magnesia rod is calcined shortly, cooled and dipped into Zn(OH)2. This procedure is
repeated until some white ZnO sticks to the tip of the rod. The tip is then dipped into 1%
(weight) Co(NO3)2 solution and calcined thoroughly. Besides a blackening caused by cobalt
oxide traces of the green zinc-cobalt mixed oxide should appear (Rinmann’s Green).
-
Detection of Ca2+: the remaining solution of the cation separation sequence is reduced in
volume to 4-5 ml by evaporation in beaker. If any crystallization starts, the evaporation
must be stopped, however, and water is added until everything is just dissolved again. A
spatula tip quantity of NH4Cl is dissolved in half of the solution and K4[Fe(CN)6] solution is
added. A light-colored precipitate, (NH4)2Ca3[Fe(CN)6]2, indicates the presence of Ca2+.
The flame color of Ca2+ can only be observed through a spectroscope or a yellow filter,
since large quantities of sodium were introduced by the reagents added.
-
Anions: the solution remaining after the carbonate precipitation is now examined. For that
purpose it is acidified first with sulphuric acid, until the pH drops below 2. Attention:
violent CO2 evolution!
-
Iodide: a small amount of sodium or potassium nitrite is added to the acidified solution. A
dark color indicates iodine. The solution is covered with a layer of hexane or light
petroleum and shaken. The hydrocarbon layer becomes purple because of the iodine. The
hydrocarbon is cautiously removed by sucking it off with Pasteur pipette. Fresh
hydrocarbon is added and the mixture is shaken again thoroughly, the hydrocarbon is
removed again. This extraction cycle is repeated until the hydrocarbon remains clear. Test
whether the pH is still acidic and add H2SO4 if this no more the case. Add also some more
nitrite. If iodine is formed again, the extraction, acidification and nitrite addition are
97
repeated over and over, until no iodide remains in the sample. It is very important to
complete the iodide removal because of interferences in subsequent operations.
-
Bromide: The solution remaining after complete iodide removal is boiled in a ventilated
hood until the evolution of brown vapors (NO2• from the decomposition of nitrite) ceases.
Now some crystals of potassium permanganate are added, and a fresh layer of hydrocarbon
is placed on top. By shaking the permanganate is dissolved and it oxidizes bromide to
bromine, Br2, which appears as a yellow to brown color in the hydrocarbon. The solution
may become dark because of permanganate reduction products. Now the bromide has to be
removed completely in a similar cycle as with the iodide. The hydrocarbon layer is
replaced, pH is controlled and permanganate is added repeatedly until the hydrocarbon
remains clear. The solution should be acidic and dark purple because of non reacted
permanganate at the end.
-
Chloride: the purple permanganate is reduced to colorless Mn2+ by cautious addition of
hydrogen peroxide. The solution is boiled shortly to eliminate excessive H2O2. Silver nitrate
solution is added and white AgCl should precipitate. The precipitate is centrifuged and
washed 2 times with 5 ml of water. The washing solutions are discarded. 5 ml of H2O and
1 ml conc. NH3 are added. Upon shaking, the precipitate should begin to dissolve under
formation of Ag(NH3)2]+. AgI and AgBr are not soluble under this condition and show
yellow colors.
98
Separation scheme of the exercise in qualitative analysis
Start solution, weakly acidic
CO32Precipitate
Solution
FeCO3, ZnCO3, CaCO3
I-, Br-, ClH2SO4
CH3COOH
Solution
Solution
Fe2+/3+, Zn2+,Ca2+
I-, Br-, ClNO2-
NH3
Precipitate
Solution
Solution
Extract
Fe(OH)2/3
Zn(NH3)42+, Ca2+
Br-, Cl-
I2
MnO4-
-NH3, pH=8-9
H2O2
Precipitate
Precipitate
Solution
Solution
Extract
Fe(OH)3
Zn(OH)2
Ca2+
Cl-
Br2
Detection
HCl
Ag+
Solution
Solution
Precipitate
Fe3+
Zn2+
AgCl
Detection
Detection
Detection
HCl
99
XI Mineral synthesis and analysis
Synthesis and analysis of struvite Mg(NH4)(PO4)6H2O
Molar mass: 245.45 g/mol
Struvite is a biogenic mineral which is formed from excrements of sea birds if the material is
exposed to only little precipitations and bacterial decomposition. It is the major constituent of
Chilean guano and used as a powerful fertiliser. Struvite forms nice crystals and can be analysed
easily for its components.
Preparation:
Dissolve 1.74 g of K2PO4 (not KH2PO4!!!) and 0.54 g NH4Cl in 20 ml water in a beaker. It is
crucial to use potassium salts, the corresponding sodium compounds are not soluble enough! Add
0.75 ml of concentrated NH3 (ventilated hood, cover beaker with a watch glass!). In another beaker
either 0.95 g MgCl2 or 2 g MgCl26H2O or 1.85 g Mg(NO3)22H2O or 2.56 g Mg(NO3)26H2O are
dissolved in 10 ml H2O.
The phosphate/ammonium mixture is stirred vigorously and magnesium solution is added
dropwise. As soon as the first turbidity appears the addition is stopped for 2 minutes and then
continued slowly until all magnesium is consumed. A drop of the suspension can be transferred
onto a slide and is observed under a microscope (80-200 fold magnification). The reaction mixture
is finally cooled for 15 minutes on an ice bath. Filter with vacuum assistance on a glass filter frit
and wash first with an ice-cooled mixture of 0.4 ml concentrated NH3 and 50 ml H2O, then with
little ice-cooled water.
Analyses:
Magnesium, ammonium and phosphate are all possible.
Mg2+ can be determined by complexometric analysis. However, phosphate has to be removed first
since it precipitates Mg2+ at the pH required for the EDTA titration. This is done as follows:
dissolve a weighed sample of struvite in dilute hydrochloric acid (phosphate buffer is formed), but
100
do not allow the solution to become too acidic (pH = 2-3). Add 0.1 M Fe(NO3)3 until the phosphate
precipitation ends and the Fe(OH)3 precipitation starts. Add the ammonia buffer required for the
complexometric titration. Excess Fe(III) is precipitated completely as Fe(OH)3. Filter through a
folded paper filter and wash with ammonia buffer in order to transfer the Mg2+ quantitatively into
the receiving flask (conical flask). Add Erio T indicator and titrate with EDTA solution of known
concentration.
Ammonium can be isolated and determined by Kjeldahl distillation. Prepare 100 ml 0.1 M boric
acid solution and set up a distillation apparatus. A sample of struvite is weighed into the distillation
flask (50 ml) and 15 ml 1 M NaOH is added. The flask is immediately mounted. The receiving
flask is filled with 20 ml of boric acid solution and a stirrer bar is added. This receiving flask also
immediately attached with a vacuum adapter. A stirrer is placed under the receiving flask and
activated. Heat the sodium hydroxide to boiling and keep this for 20 minutes. Afterwards the boric
acid solution is titrated with hydrochloric acid of exactly known concentration, with methyl red or
methyl orange as indicator. Consider which reactions take place in the distillation apparatus and
during the titration.
101
Appendix
pk values of some acids at 25°C
+
Ammonium (NH4 )..............................9.23 .................................. 0 M
Boric acid (B(OH)3)............................9.14, 12.74, 13.80............. 0.1 - 0.01 N
Acetic acid (HOAc, CH3COOH).........4.75
‘’
Ethylenediamine tetraacetic acid .......2.0, 2.69, 6.18, 10.15........ 0.1M KNO3
(H4EDTA, H4Y)
+
+
Glycinium (Hgly , NH3 CH2COOH) ...2.35, 9.78 ......................... 0.1 - 0.01 N
Carbonic acid (H2CO3) ......................6.37, 10.25
‘’
Oxalic acid (H2ox, HOOCCOOH) ......1.23, 4.19
‘’
Phosphoric acid (H3PO4)...................2.12, 7.21, 12.67
‘’
Nitric acid (HNO3) ..............................-1.43 ................................. 0 M
Hydrochloric acid (HCl)......................-6.1 ................................... 0 M
Sulphuric acid (H2SO4)......................-8.0, 1.92 .......................... 0 M
Hydrogen suphide (H2S)....................7.04, 19 ............................ 0.1 - 0.01 N
+
+
"Htris " (NH3 C(CH2OH)3) .................8.09 .................................. 0.1 M KNO3
102
pH indicators
Indicator
Alizarine yellow
Bromocresol green
Bromocresol
purple
Bromothymol blue
Cresol red
Litmus
Methyl red
Neutral red
Phenolphthalein
Pheno red
Thymol blue
Acidic colour
yellow
yellow
yellow
yellow
yellow
yellow
red
red
red
colourless
yellow
red
yellow
pH range of colour
change
10.4 - 12.0
3.8 - 5.4
5.2 - 6.8
6.0
0.4
7.0
4.4
4.8
6.8
8.2
6.6
1.2
8.0
- 7.6
- 1.8
- 8.8
- 6.2
- 6.0
- 8.0
- 10.0
- 8.0
- 2.8
- 9.6
pK
11
4
5
Alkaline colour
pink
blue
violet
7
8
6
5
7
9
7
2
9
blue
red
violet
blue
yellow
yellow
red
violet
yellow
blue
Further the following indicators are used:
for complexometry:
Murexide
Erio T
acidic colour
violet
violet-brown
basic colour
blue
blue
and indicator paper. This is a mixed indicator, which shows colour variations over a wide pH
range. It is composed of methyl red, dimethylamino azobenzene, bromothymol blue and thymol
blue, the single components being present in 0.025-0.1% concentrations. Die colour shown
depending on the pH value is: pH 3 - red - orange - yellow - green - blue - pH 10.
103
Standard reduction potentials
0
E (V)
+
Ag /Ag ............................................................0.7996
AgCl/Ag ..........................................................0.2223
Br2/2 Br ..........................................................1.065
4+
3+
Ce /Ce (1 M H2SO4)................................1.4587
Cl2/2 Cl ..........................................................1.3583
3+
2+
Co /Co (3 M HNO3) .................................1.842
3+
2+
[Co(NH3)6] /[Co(NH3)6] ...........................0.1
3+
2+
Cr /Cr ....................................................... -0.41
23+
Cr2O7 /2 Cr ................................................1.33
+
Cs /Cs ........................................................... -2.923
+
Cu /Cu ............................................................0.522
2+
+
Cu /Cu .........................................................0.158
2+
Cu /Cu...........................................................0.3402
3+
2+
Fe /Fe ........................................................0.770
34[Fe(CN)6] /Fe(CN)6] (1 M H2SO4) ..........0.69
+
2 H /H2 ............................................................0.0000
H2O2/2 H2O ...................................................1.776
Hg2Cl2/2 Hg (Kalomel) (satd. KCl) .............0.2415
I2/2 I ...............................................................0.535
+
K /K............................................................... -2.924
+
Li /Li ............................................................. -3.045
2+
MnO2/Mn .....................................................1.208
2MnO4 /MnO4 ...............................................0.564
-
MnO4 /MnO2 ..................................................1.679
-
2+
MnO4 /Mn ....................................................1.496
+
Na /Na .......................................................... -2.7109
2+
Ni /Ni........................................................... -0.23
O2/2 H2O ........................................................1.229
+
Rb /Rb .......................................................... -2.925
2S/S ............................................................... -0.508
2-
2-
S4O6 /2 S2O3 .............................................0.09
2+
Zn /Zn ......................................................... -0.7628
104
Complex formation constants
a) EDTA-complexes (pK values of H4Y: 2.0, 2.69, 6.18, 10.15; 20°C, 0.1 M KNO3)
M
T (°C) Medium
log K1
2+
20
0.1 M KNO3
Mg
2+
Ca
25
‘’
2+
Sr
20
‘’
2+
Ba
25
‘’
3+
20
‘’
Al
3+
Sc
22
0.5 M NaCl
3+
La
20
0.1 M KNO3
2+
20
‘’
Mn
2+
Fe
20
‘’
2+
Co
20
‘’
2+
Ni
20
‘’
2+
Cu
20
‘’
2+
Zn
20
‘’
2+
Cd
20
‘’
2+
Hg
25
‘’
2+
Pb
20
‘’
+
b) NH3 complexes (pK of NH4 : 9.24, 25°C, I
M
T (°C) Medium
2+
Co
2+
Ni
2+
Cu
2+
Zn
+
Ag
30
‘’
‘’
‘’
25
2 M NH4NO3
‘’
‘’
‘’
‘’
5.6
4.72
4.2
3.8
14.8
21.84
13.0
11.7
10.7
13.8
17.4
19.6
13.1
10.4
17.50
13.5
0)
log K1 log K2 log K3 log K4 log K5 log K6
2.11
2.78
4.14
2.45
3.35
1.63
2.27
3.52
2.28
3.90
1.05
1.65
2.87
2.64
0.76
1.31
2.15
2.11
0.18
0.65
-0.62
0.08
c) Glycine complexes
+
+
pK values of H3 N-CH2-COOH (Hgly ): 2.35, 9.78 (25°C, 0.1 M KNO3)
M
T (°C) Medium
2+
Cu
2+
Ni
25
‘’
0.1 M KNO3
‘’
log K1 log K2 log K3
8.23
5.73
6.96
4.83
3.44
105
Solubility products
MaLb(s)
a M+ b L
a
KL = [M] .[L]
b
Solid
pKL (at 25°C)
AgCl
AgBr
AgI
Ag2S
CdS
CuS
FeS
HgS
MnS
NiS
PbS
ZnS
9.80
12.27
16.07
50.22
25.10
36.22
18.22
52.70
13.52
20.97
27.52
24.70
106
Conductivity data
-1
2
Molar asymptotic conductivity 0 of ions in water at 25°C (in Scm mol )
+
+
H ............349.65
+
Na ............50.08
+
K ..............73.47
3+
Ag ..............61.9
2+............
Ba
127.2
2+
Co ............110
-
La ............. 209.1
+
N(C2H5)4 .......32.6
3+
Co(NH3)6 . 305.7
-
OH ...............198
-
2-
NO3 ..........71.42
SO4 .............. 160
C2O4 ...148.22
ClO4 ................67.3
CO3 .......138.6
Fe(CN)6 .... 302.7
2-
Cl ..............76.31
-
-
2-
Br ................78.1
3-
Temperature dependence of 0: example HCI

T (°C)
-1
2
0 (Scm mol )
5
296.4
15
360.8
25
424.5
35
487.0
45
547.9
55
606.6
Concentration dependence of : example NaCI
-1
c (moll )

-2
0
-1
 (Scm mol )
126.39
-4
-3
-3
-2
-1
5.10
10
5.10
10
10
124.44
123.68
120.59
118.45
106.69
2
-1
Some values of 0 in non-aqueous Media at 25°C (Scm mol )
Solvent
CH3OH
C2H5OH
HCN (18°C)
NaCI KCI
98 105
42
45
363 363
KBr
109
Dielectric constant
32.6
24.3
118
107