Download General Chemistry - Bioinorganic and Solution Chemistry Group

General Chemistry
A Practical Course for Students of
Pharmaceutical Sciences and Biology
ETH Zurich 535-1001-00
2016
1
Contents
Introduction............................................................................................................................................. 5
Glass ...................................................................................................................................................... 10
Properties of glassware ..................................................................................................................... 10
Cleaning of glassware ........................................................................................................................ 11
Glass figuration.................................................................................................................................. 12
Glass cutting .................................................................................................................................. 12
Smoothing ..................................................................................................................................... 12
Tube bending ................................................................................................................................. 12
Volumes of laboratory glassware ...................................................................................................... 13
Is volume a conserved quantity (do volumes add up)? .................................................................... 15
Fractioning Methods I ........................................................................................................................... 17
Precipitation of a coagulate: Fe(OH)3 ................................................................................................ 17
Precipitation of calcium carbonate CaCO3, filtration with a porcelain Buchner funnel.................... 18
Synthesis of hydroxyapatite; filtration with glass frit crucible, pore size 4 .......................................... 19
Precipitation of AgCl and Ag2CrO4 ..................................................................................................... 22
Precipitation of colloidal (nanoparticulate) Prussian Blue ................................................................ 22
Preparation:....................................................................................................................................... 22
Re-crystallisation of a mixture of KNO3 with Cu(NO3)2: a purification method ................................ 23
Crystalline solids .................................................................................................................................... 25
Solid mixtures – mixed crystals ......................................................................................................... 26
Mixed crystals .................................................................................................................................... 26
Solubility of NH4Cl, KNO3, solubility product of KClO4....................................................................... 27
Argentometric titration ......................................................................................................................... 31
End point indication of some argentometric titrations..................................................................... 33
Method of Fajans........................................................................................................................... 34
Preparation of 0.05 M AgNO3 solution and its calibration ............................................................ 34
Argentometric titration of Br-, I-, SCN- ........................................................................................... 35
Fractioning 2 .......................................................................................................................................... 36
Condensation .................................................................................................................................... 36
Sublimation ....................................................................................................................................... 37
Distillation.......................................................................................................................................... 37
Vacuum distillation ............................................................................................................................ 38
2
Distillation of an azeotropic two-component mixture ...................................................................... 39
Volatile Compounds .............................................................................................................................. 42
Determination of melting and boiling points ........................................................................................ 42
Preparation of volatile substances .................................................................................................... 44
Determination of the molar mass by melting point depression (cryoscopy) ....................................... 45
Acids and Bases ..................................................................................................................................... 47
Acid and base definitions by Lewis and Brønsted-Lowry .................................................................. 47
Proton transfer in aqueous solution ..................................................................................................... 49
pK values, pH concept, strong acids and bases, weak acids and bases, multistage deprotonation . 49
Strong acids ........................................................................................................................................... 50
Weak acids............................................................................................................................................. 50
Concept of pH ........................................................................................................................................ 51
Solutions with fixed pH: buffers ............................................................................................................ 51
pH measurement ................................................................................................................................... 52
Synthesis of ethanolic hydrogen chloride ............................................................................................. 52
Esterification of boric acid ..................................................................................................................... 54
Reaction between gaseous NH3 and HCl ............................................................................................... 54
Synthesis of a calcium salt ..................................................................................................................... 55
Sublimation of ammonium chloride...................................................................................................... 55
NH3/HCl ............................................................................................................................................. 55
Aluminium chloride as Lewis acid ......................................................................................................... 55
Preparation of AlCl3. Reaction of AlCl3 with ether. Reaction of AlCl3 with KCl.................................. 55
Preparation of potassium hydrogen tartrate ........................................................................................ 57
Water as acid, water as base ................................................................................................................. 58
Acidimetric titration .............................................................................................................................. 59
Determination of the pK values of the indicator thymol blue .............................................................. 60
Preparation of a phosphate buffer of pH = 7.30 and I = 0.16 ............................................................... 62
Redox Reactions .................................................................................................................................... 65
Redox reactions, solvent free or in aqueous solutions ..................................................................... 65
Thermal decomposition of potassium chlorate .................................................................................... 66
Preparation of CuCl ............................................................................................................................... 68
Redox reactions in qualitative analysis ................................................................................................. 68
Detection of chromium as chromate. ................................................................................................... 69
Proof of oxidising agents. Conversion of I- into I2.................................................................................. 69
3
Disproportionation of H2O2, catalase .................................................................................................... 70
Standard reduction potential Fe(CN)63- / Fe(CN)64- ............................................................................. 70
Permanganometric titration.................................................................................................................. 72
Permanganometric determination of oxalic acid, (COOH)2 .................................................................. 73
Iodometric titration (of Cu2+ solution)................................................................................................... 73
Ligand Exchange and Complex Formation ............................................................................................ 75
Introductory experiments in coordination chemistry ........................................................................... 79
e) Inertness of Fe(CN)64- ........................................................................................................................ 82
Preparative coordination chemistry...................................................................................................... 83
Tetraammine nickel nitrite Ni(NH3)4(NO2)2 ....................................................................................... 84
Potassium dioxalato cuprate(II) K2Cu(OOCCOO)22H2O ................................................................... 85
Metal indicators .................................................................................................................................... 86
Dithizone as metal indicator ................................................................................................................. 87
Determination of the hardness of water by complexometric titration ................................................ 87
Chromatography ................................................................................................................................... 90
Chromatographic separation of dyes .................................................................................................... 90
Liquid-liquid distribution ....................................................................................................................... 91
Determination of the distribution coefficient of iodine in the solvent system H2O/CH2Cl2 ................. 93
Ion exchangers ...................................................................................................................................... 93
Ion exchange chromatography: separation of Cu2+, Ni2+, Fe3+ .............................................................. 95
Qualitative analysis................................................................................................................................ 97
Appendix.............................................................................................................................................. 104
pk values of some acids at 25°C ...................................................................................................... 104
Standard reduction potentials ......................................................................................................... 106
Complex formation constants ......................................................................................................... 107
Solubility products ........................................................................................................................... 109
Conductivity data ............................................................................................................................ 110
4
Introduction
Objectives
The beginner’s course in inorganic and general chemistry is intended to teach the fundamental
methods of chemical laboratory work to students of biology and pharmaceutical sciences and to
make them familiar with the important reaction types in inorganic chemistry. Restrictions concerning
the availability of instrumentation, laboratory space, personnel and time call for a certain degree of
flexibility. Inorganic chemistry, despite its name, is essential for biological systems. All organisms
absorb inorganic nutrients, which are used then to fulfil very special functions, like formation of
skeletons and shells, reaction centres in biological catalysts (enzymes) and transport mechanisms
(e.g. oxygen transport).
The course presents a variety of materials to the student. Quantitative analyses, some of them
carried out with instruments, require accurate and clean working skills. It is expected that the
students exert themselves to acquire these in the very beginning, because according to experience, it
is difficult to improve later during advanced courses. You should try to understand why operations
are carried out in a certain way, and what the consequences of alternative approaches would be. In
case of doubt you should always consult a teaching assistant if a method does not make sense to
you.
The laboratory course and the first-semester lecture have some common contents, albeit not very
extensively, because the initial part of the lecture cannot be reproduced in the laboratories we use.
Therefore, the course manual contains theoretical considerations which precede the contents of the
lecture. It is recommended to study these paragraphs before doing the corresponding experiment
and to complete by information found in books. 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.
5
Safety
Laboratory work can be dangerous, like any kind of practical work. The worst hazards threaten the
eyes. Therefore, it is indispensable to wear
SAFETY GLASSES
This is permanently necessary, because you do not know what your neighbours are carrying out at
the moment. It is important not only to protect yourself from your own misfortune, but also from the
whole environment. These hints do not mean that you are working in an extremely dangerous area!
It is more likely to have an accident in your kitchen at home because of the sloppier precautions
taken. Fear is no good advisor, also in a chemistry laboratory. The best prevention is knowledge. As
long as you are sure about what the properties of the material in front of you are, you are safe. A
kilogram of potassium cyanide on your bench poses no risk, as long as you do not try to eat from it or
come up with the idea to add acid which liberates hydrogen cyanide gas. From these considerations
the most important statement of chemical safety follows:
-
Solids and liquids are only a problem when you ingest them. This may happen unconsciously,
e.g. by contaminated hands. There is currently a strong impulse to wear gloves even during
the most harmless operations in the laboratory. Unfortunately in most cases this is not a
solution but rather a relocation of the problem. Whenever you touch door handles or other
public installations with contaminated gloves, you are protected until you remove them ...
Question: how to remove a pair of contaminated gloves without touching the
contamination? (Wash them, but you have to touch the spigot first?!) You also endanger
your colleagues who might not be aware that you just touched door handles with dirty
gloves. It is better to work without gloves except when using very aggressive substances.
Instead, wash your hands immediately if you suspect contamination. Most (but not all) react
very slowly with human skin.
-
Gases, if toxic, are a real problem because they can reach you without your assistance. The
only safe working area in this case is the fume hood. This is a fairly safe approach as long as
the ventilation is running properly and you do not poke your head into the hood.
6
Special risks:
-
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 gas burner even some meters away.
-
Poisonous chemical waste is not to be disposed of into the sink but in special 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).
At the very start of the course the proper methods to use the equipment are introduced by the
teaching assistants. It should not be difficult for the students to work safely and cleanly after this.
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 gas 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 by 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 sideways. Never shake up and down, or direct the opening of a test
tube towards a person! 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 can be placed on a socket (e.g. “Labor-Boy” in order to make contact. In case
of overheating the socket can be lowered or removed instantly.
7
The gas 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
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 is, the faster the reactions are. The typical concentration unit is mol/l,
abbreviated M. Unfortunately, manufacturers of chemicals indicate concentrations in solution in
8
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
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. At the beginning of the
course you will obtain some exercises referring to this.
9
Glass
In this first section of the course the students are acquainted with properties of the most important
basic material of the chemistry laboratory. Glass is different in many respects from stone, metals,
wood and plastics, the other abundant basic materials. Despite its brittleness and its tendency to
form dangerous splinters it has unsurpassed qualities for the handling of substances. It is hardly or
not at all attacked by most chemicals, it supports strong heating and cooling, albeit not suddenly, and
it is transparent such that events in a vessel can easily be observed.
Further, we introduce a dosing device that has become popular in life sciences: the piston-driven air
displacement pipette.
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 gas burner.
Abbreviations for glassware used in this manual (German version)
RG: test tube
BG: beaker
Properties of glassware
Pressure resistance
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. It is even more dangerous to apply pressure to glass
containers. While under vacuum the pressure difference across the wall of the vessel cannot be
greater than one atmosphere, it can be almost unlimited on connection to sources of compressed
gases. Furthermore, splinters will fly away outward, other than in case of implosion.
10
Heating/cooling resistance
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. Glass is a poor
heat conductor; therefore a thick wall expands faster on the heated outer side than on the inside.
The resulting tension breaks the glass. Quartz, despite its similar properties, resists sudden heating
and cooling, because it does hardly expand or contract. However, it is difficult to process and
correspondingly expensive.
Precision of graduations on glassware
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 volume) and pipettes (calibrated on release
volume).
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.
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.
11
Glass figuration
Every student of sciences should be familiar with the most simple glassblowing operations. 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. The pieces are kept for later use.
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.
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 and let the tube to bend in the flame by its own weight.
Never force bending, instead increase heating.
12
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.
d = 2r
h
r
r
Temperature ° C
Density of water (g/ml)
10
0.999700
15
0.999099
20
0.998203
25
0.997044
30
0.995646
13
The tare values of the graduated cylinders in the inventory are determined. Then, 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 with the density of water and compare with the
indications.
Exact volume of the 100 ml graduated flask
The tare of a dry 100 ml graduated flask is measured, and 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. The procedure is repeated
again, 8 times.
The mean and the median of the 9 measurements of the filled-in volume are determined. The
median is the middle value if you list all of them ordered by magnitude. In our case this is position 5.
We test now whether the mean (average) lies in the 75% confidence interval around the predicted
value (100 ml). This is done by the single sample bilateral t-test. We must calculate the according tvalue for our 9 measurements (this is the single sample) und compare to tabulated values of the tfunction. The t-value is determined as follows: t  n
V  V0
, where n=9 represents the sample size,
s
V the mean, V0 the expected value and s the standard deviation. s is defined as s 

n
i 1
(Vi  V )2
n 1
and available on pocket calculators. The t-value is compared to the tabulated number of the 75%
confidence interval of the t-distribution.For 8 degrees of freedom (9-1, because the first
measurement does not allow for comparison) this is under bilateral condition (mean can be larger or
smaller than the expected value) t0.75=1.240. If |t| > t0.75, the flask is calibrated incorrectly, given 25%
probability of error.
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.
14
Determine the total weight of the bottle and calculate the drained volume. Take 9 measurements in
total and analyse as you did with the graduated flask.
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 = π 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?
Piston pipettes
Piston-driven air displacement pipettes are not glassware, but they replaced some in biology and
biochemistry laboratories. Their advantage is the rapid dosage of small volumes, even with different
amounts and solutions. This is achieved by interchangeable pipette tips made from plastic. In a
normal chemistry laboratory, the usage is limited because piston pipettes are sensitive against
15
corrosive chemicals, and not precise with volatile solvents. Operation is simple: the pipette body
contains a hollow cylinder with a piston that travels an exactly known distance, when it is moved by
the built-in mechanics. The cylinder is coupled to the pipette tip by an air duct. The tip is just plugged
onto the outlet of the duct. The mechanics has three notches: a stop, where the mechanism is in
resting position, a threshold resistance at the dosage volume, and another stop at complete
expulsion. The piston is held by an expanded spring in the resting position.
Usage is easy, but has to be exercised for reliable work. Since many course participants will be
working in biologically oriented laboratories in the future, we teach the handling already here. We
use a mid-size version that can dose 0.1 – 1 ml. The entire volume range covered by piston pipettes is
0.001 – 10 ml. Operation:
-
Set volume, if the pipette is the variable variant. Here the greatest differences are found
among the various brands.
-
Hold the pipette in your preferred hand. Mount a pipette tip, air-tight but not too tough.
-
Put thumb on the plunger button and push down to the volume threshold.
-
Sink pipette tip into the solution and let the plunger come back slowly to the resting position
under control of the thumb. Liquid is aspired into the tip.
-
Never let the plunger shoot up! This would draw some solution into the interior cylinder!
-
Withdraw pipette from solution, direct tip into target and, optionally, attach to container
wall.
-
Push plunger button forcefully to the total expulsion stop, beyond the volume threshold.
-
Return plunger to resting position.
If the same solution is dosed repeatedly, the tip can be retained; the volume can be varied, of course.
For dosage of a different solution the tip is removed by means of the built-in mechanics, and a new
one is mounted. Used tips are combustible waste.
Exercises with a piston pipette of 0.1 - 1.0 ml
A simple way to test your own skills with a piston pipette is to weigh doses of a liquid with known
density. It is required that your weighing skills are developed. You need:
Piston pipette and suitable tips
Balance with at least 3-digit display
Two beakers of 50 ml
16
Water and density table thereof
You can find a density table above, under “Volume determination of test tubes and graduated
cylinders”.
One beaker is half filled with water and serves as reservoir. In contrast to normal glass pipettes, the
extraction of liquid from narrow-necked containers is often difficult with piston pipettes. The other
beaker is placed on the balance, which is set then to zero. A selected volume is dosed from the
reservoir into the beaker on the balance. The weight is read quickly before evaporation causes
losses. The balance is set to zero again, and the procedure is repeated. We do this exercise 9 time
with 1 ml and 9 times with 0.1 ml.
Both series are analysed by the same statistical method as before. What follows from the result?
Fractioning Methods I
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 a coagulate: Fe(OH)3
Filtration with filter paper
Folded filters serve for the quick filtration of small amounts of precipitates or impurities like lint etc.
from large volumes of liquid. Precipitate and filter are discarded.
17
Coagulates are non-crystalline precipitates and often only approximately stoichiometric. They may
contain variable amounts of solvent which makes them jelly-like. They can be by-products of desired
reactions and are filtered with a normal filter paper in a funnel, under gravity. 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 2 M 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.
Reaction:
dissolved
Fe3  3OH  

 Fe(OH )3
coagulated
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. When the flask is emptied, its interior is rinsed from top
to the bottom. The collected solution added in the funnel to the previously collected material. For
the most accurate analytical work this is repeated two times more. Finally, the collected precipitate is
washed cautiously with some deionised water by rinsing it downward, from the upper lip of the
funnel. This kind of flushing is necessary for exact and quantitative work.
Precipitation of calcium carbonate CaCO3, filtration with a porcelain Buchner funnel
Calcium forms, together with CO2, a sparingly soluble ionic compound under neutral to alkaline
conditions: CaCO3, calcium carbonate, also called limestone. It is remarkable that most the calcium
carbonate on Earth is a product of life! Exoskeletons of molluscs, corals and crustaceans contain
biogenic calcium carbonate. Fossil organisms formed most of the limestone found today. We prepare
calcium carbonate in the laboratory as follows:
Reaction:
dissolved
Ca 2  CO32 

 CaCO3
crystalline
18
Preparation:
-
In a 100 ml beaker 2 g of calcium chloride CaCl22H2O are dissolved in 20 ml 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 rubber seal is placed into the neck of a suitable suction bottle (Buchner flask, thickwalled!) 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 substantial
repair cost.
The contents of the 50 ml beaker are poured slowly under stirring into the calcium chloride solution.
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 attached to 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! The filter paper with the cake can be separated from the
Buchner funnel by means of a spatula.
Synthesis of hydroxyapatite; filtration with glass frit crucible, pore size 4
Introduction:
Hydroxyapatite is the mineral on which bones and teeth are based on. It consists of calcium ions,
phosphate ions and hydroxide ions: Ca5(PO4)3(OH). The hydroxide ions can be partially substituted by
fluoride ions (tooth enamel; see also experiments with mixed ions in the following chapter). In the
biological environment crystals of the material are grown by specialized cells. Their control of
19
concentrations produces well-defined shapes and sizes even at ambient temperature. In the
synthesis under laboratory conditions, the crystallization is too rapid at room temperature, the
starting material is consumed in short time and the crystals remain small, which makes filtration
almost impossible. At elevated temperature, crystallization becomes slower with increasing
solubility. Crystallization from hot solutions followed by slow cooling usually leads too reasonable
products. Since, depending on the composition of the staring solution, calcium phosphates of
different stoichiometry (Ca3(PO4)2, CaHPO4) can be formed, it is indispensable to dose the reagents
precisely.
Reagents:

0.370 g Ca(OH)2, weigh as accurately as possible

6 ml 0.50 M H3PO4

Ethanol and acetone
Procedure:
0.37 g Ca(OH)2 are weighed directly and as precisely as possible into a 100 ml beaker. Add 50 ml of
deionized water and a magnetic stirrer bar, and heat to boiling under slow stirring on the electric
heater/stirrer. Under continued boiling and stirring, add 6 ml 0.50 M H3PO4 in increments of 1 ml
every minute. Test the pH with universal indicator paper; it should lie between 6 and 7. If the paper
turns deep green to blue, add H3PO4 drop wise until the solution appears neutral. The solution is kept
boiling for another 30 minutes and then allowed too cool to room temperature. Set up the vacuum
filtration apparatus as depicted below on the right, with the filtration crucible type G4. The filtration
is carried out as before with the Buchner funnel, the precipitate is stirred up and passed along a glass
rod into the crucible. When all material has collected in the crucible, 200 ml of deionized water are
passed through. The crucible is separated from the suction flask and the solution is discarded. The
crucible is placed again on the flask, and we pass 50 ml ethanol through the precipitate, followed by
50 ml acetone. The crucible is transferred into a desiccator and vacuum is applied for 1 hour. During
the boiling and the drying, another experiment should be started in parallel. After the vacuum drying,
the white powder is transferred into a pill tube by means of a spatula. The pill tube is labeled and the
hydroxyapatite is kept for an experiment in the following chapter.
20
After filtration, some material always remains stuck in the crucible. Hydroxyapatite is easily removed.
We mount the crucible in inverted position on the suction flask (depicted on the left side), apply
vacuum and pass through a few ml of 1 M HCl. This dissolves the hydroxyapatite, and after flushing
with deionized water the crucible should be clean. There are precipitates much more resistant than
our example; however, the operative principle remains the same: passing solvent in reverse direction
dissolves best materials that have become stuck.
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 retarded by hand only very gently; otherwise the Coriolis
forces will transfer rotational energy to the liquid and cause the precipitates to whirl up again.
21
Precipitation of AgCl and 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. Furthermore, they
can be used for quantitative determinations because of the extremely low solubility of the
compounds. Silver are toxic to microbes, therefore, soluble silver compounds, especially the nitrate,
were used for disinfection purpose. Today they are used not so frequently, mainly because of
aesthetic reasons: since Ag+ is easily reduced to elemental Ag0, the disinfected zone on the skin
acquires a dark hue. It takes weeks until the harmless colouration disappears. Because of this effect,
silver nitrate was also called “lunar caustic” in the past.
Selective precipitation and centrifugation of Cl- and CrO42- as silver salts
About 2 ml of the supplied solution mixture of 0.1 M sodium chloride NaCl and 0.1 M potassium
chromate K2CrO4 are transferred into a small test tube and 0.2 M silver nitrate solution (AgNO3) is
added drop wise, 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
Preparation:
0.4 g potassium hexacyanoferrate(II) K4[Fe(CN)6] are dissolved in 100 ml water
0.3 g iron(III) chloride FeCl3 6H2O in 50 ml water
22
The FeCl3 solution is added to the K4[Fe(CN)6] solution 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 < 10-3 mm 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 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 lint
and dust particles.
Colloidal suspensions have much in common with suspensions of single-cell organisms, even with
“solutions” of proteins and other macromolecules. Liquid nanoparticles suspension is a “modern”
expression for colloid. However, nanoparticles can also be suspended in gases or be adsorbed on
surfaces.
Re-crystallisation of a mixture of KNO3 with Cu(NO3)2: a purification method
Re-crystallisation is one of the most important purification techniques in chemistry, probably the
most important in organic chemistry. It is founded on the concepts that the solubility of different
solids does hardly change to the same extent with temperature, and that growing crystals tend to
incorporate preferably molecules or ions they already consist of. In general, the method is carried
out as follows: a saturated solution is made up at elevated temperature, usually at boiling point of
the solvent. This is achieved by adding a small amount of solvent to the solid, followed by warming
and slow further addition of solvent under continued warming, until all solid is just dissolved. Then,
the solution is allowed to cool slowly, and the least soluble component will separate as rather pure
crystals. Depending on the special problem the solution can be cooled further, close to the melting
point of the solvent, in order to improve the yield. This approach can be counterproductive because
impurities could start to co-crystallize. However, minority components tend to remain in solution
because of their low concentrations. The crystals are collected by one of the filtration techniques
mentioned before, and they are rinsed with pre-cooled pure solvent in order to remove the adherent
stock solution. This step always causes some loss. The whole procedure can be repeated with the
purified crystals over and over to obtain a cleaner product every time, though the losses limit the
method.
23
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 3H2O 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.
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.
24
Crystalline solids
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 Xray 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.
Solubility: a key problem in biology
Many nutrients are abundant in nature, but often not easily available to organisms. This is frequently
caused by the poor solubility of the compounds, salts or complexes that contain the material. An
example is iron, an important nutrient for all kinds of life (active sites in enzymes, haemoglobin ...).
Because of atmospheric oxygen iron in contact with air exists predominantly as Fe3+. This ion forms
almost insoluble Fe(O)OH (approximate composition) at pH 7 in water. According to calculations,
about 109 Fe3+ ions per litre remain in solution. This is next to nothing compared to typical
concentrations of functional small molecules in organisms, which are 1015 to 1022 ions per litre. Life
has developed numerous strategies to safeguard iron supply.
Another example is the solubility of Ca2+, also important for all organisms. Some Ca2+ salts are rather
soluble, while others are dissolved only to a limited extent. In order to build up bone or dental
enamel, Ca2+ is absorbed in soluble form, and can be found in extracellular fluid in concentrations of
1-2 mM. Calcium in blood consists of 50% hydrated ions, 35% are protein bound (albumin and
globulines), and 15% are bound by ligands (bicarbonate, lactate, citrate, phosphate). In situ it is
converted to sparingly soluble compounds under strict control. 99% of the calcium in a mammal
body are located in bones and teeth – the calcium-rich compounds hydroxyl apatite (Ca5(PO4)3(OH))
makes them stable and firm. Simultaneously, bones serve as a reservoir for calcium.
25
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. A
biological example is the above mentioned hydroxyl apatite Ca5(PO4)3(OH). The hydroxide ion of this
solid is partially substituted in dental enamel by the fluoride ion, which enhances firmness further.
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.
Mixed crystals
Mixed crystals of BaSO4 - KMnO4
Preparation:
-
In a large test tube a solution of 0.1 g K2SO4 together with little (max. 10 mg) KMnO4 is
prepared in 10 ml water.
-
In another large test tube a solution of about 50 mg BaCl22H2O in 10 ml water is prepared.
Both solutions are heated almost to boiling and the BaCl2 solution is added slowly to the mixed
solution. The precipitate is centrifuged in a centrifuge tube made from plastic. 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.
26
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.
Solubility of NH4Cl, KNO3, solubility product of KClO4
The solubility of a substance in a solvent can be expressed 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.
27
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.

 aM b  bX a 
M a X b (s) 

so: Solubility
(s): solid
K so  [M b ]a  [ X a  ]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 stoichiometric 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
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.
K so  [ K  ][ClO4 ]
for 25 ° C (density of the saturated solution d25 = 1.015 g/ml).
28
Solubility of hydroxyapatite
Introduction
The solubility of the bone mineral hydroxyapatite Ca5(PO4)3(OH), synthesized in the preceding
chapter, is even lower than that of KClO4. Parts of a skeleton should not be too soluble. In order to
detect one of the ion types in solution we need to apply a method that can measure concentrations
lower than 10–3 M. Ca2+ can be determined with high sensitivity and precision by its light emission,
when it is introduced into a hot flame. Unfortunately, the required equipment is fairly expensive and
not suitable for beginners. The hydroxide ion is a component of water, and its origin in water as a
solvent cannot be assigned. Therefore, we focus on phosphate, which can be measured by optical
detection in solution, with considerable sensitivity. For this purpose, the phosphate is converted into
a compound with molybdate:

 [(MoO3 )12 ( PO4 )]3  12H 2O
12MoO42  H3 PO4  21H  

Phosphomolybdate is intensively yellow colored. The reaction equation tells us that the compound
will form only under acidic conditions. However, we have to be aware that the yellow color may be
quenched by too much acid:
yellow

 H3[( MoO3 )12 ( PO4 )]
[(MoO3 )12 ( PO4 )]3  3H  

colorless
Therefore, the addition of acid must be sufficient, but not excessive.
Reagents:

1 g (NH4)6Mo7O24 · 4 H2O dissolved in 20 ml H2O

Phosphate standard solution (50 mg l–1 phophorus) obtained from the teaching assistant

Strip the rubber band over the glass filter crucible type G4 and insert the crucible into the
filtering receptacle

Cuvettes

Spectrophotometer
29
Procedure:
About 0.2 g of the hydroxyapatite previously synthesized are covered by 20 ml of water in a 50 ml
beaker. A magnetic stirrer bar is added, and the beaker is covered with a watch glass. The mixture is
stirred slowly at room temperature for 1 hour. During the equilibration, you should carry out another
experiment. When the equilibration period is over, the crucible in the receptacle is placed on a small
conical flask, without rubber seal and without vacuum application. Be patient, the following
procedure lasts 30-60 minutes. The slurry in the beaker is poured into the crucible and the solution is
allowed to pass slowly through the frit, driven by gravity only. The liquid filtered should be clear, not
turbid. Again, use the time to proceed with other experiments. When about 2/3 of the liquid have
passed the filter, take a 5 ml sample of the clear solution with the calibrated pipette and transfer into
a 50 ml graduated flask. 5 ml of the phosphate standard are transferred into another 50 ml
graduated flask. Add 5 ml of the molybdate solution to each of the flasks and fill both of them with
deionized water, to just below the neck. Add 0.27 ml of conc. H2SO4 to each flask, by means of a
piston pipette, and shake. Fill both flask with deionized water to the graduation mark, and
homogenize both solutions.
A 1 cm cuvette made from polystyrene (clear plastic) is filled with water and used to determine the
reference signal between 350 and 500 nm. After that, we fill the cuvette with the mixture from
phosphate standard and molybdate, and the spectrum is recorded between 350 and 500 nm. This is
repeated with the hydroxyapatite solution treated with molybdate.
Light attenuation by colored material (dyes) for a certain color (wavelength λ) is described by the law
of Lambert and Beer:
A    c  d
A: Attenuation (called absorbance or extinction)
ελ: material dependent constant
c: concentration of dye
d: thickness of the colored medium
The absorbing layer thickness in our experiments is always d = 1 cm, given by the size of the cuvette.
Also, ελ is identical at each wavelength λ, because we measure the same compound. Ergo
Astandard Asample

cstandard csample
and therefore csample 
Asample
Astandard
cstandard
30
Because we dilute the phosphorus standard by a factor of 10 in the color preparation, c1 = 5 mg l–1 of
phosphorus, corresponding to 1.6∙10–4 M phosphate. In order to obtain the solubility of
Ca5(PO4)3(OH), csample has to be multiplied by a factor of 10 because of the dilution in the color
preparation, and it has to be divided by 3, because hydroxyapatite releases 3 phosphate ions upon
dissolution. Therefore
[Ca5 PO4 3  OH] 
10  csample
3
This value should be 0.5 – 2 ∙10–3 M according to scientific reports.
Argentometric titration
Introduction
Silver ions (also Hg22+ and Tl+) instantly form stoichiometric 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). This titration method is still
the most common technique for determination of halides and is repeated in the course
“Pharmazeutische Analytik” in the 5th semester.
Titration terms:
-
Titration: step wise execution of a reaction, until a product with stoichiometric composition
is formed.
-
The equivalence or end point of a titration is reached, when the stoichiometric product is
precisely formed. Various Methods and effects are used to determine this state.
-
Receptacle: the container which normally holds the sample and where the reaction takes
place.
-
Burette: graduated tube with a stop-cock at the outlet. It is used to dose the reagent.
-
Indicator: accessory reagent for the determination of the equivalence point.
31
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 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).
The solubility product of AgCl is 10-10 M2, the one of Ag2CrO4 10-12 M3. Why is the numerical value
smaller for Ag2CrO4 though it is obviously more soluble than AgCl? Otherwise, chromate would not
be a suitable indicator.
The general reaction equation is simple:

 AgX ( s )
X   Ag  

In the titration of cyanide the process is more complex. CN- in excess over Ag+ initially forms the
colourless and soluble complex Ag(CN)2-.

 [ Ag (CN )2 ]
2CN   Ag  

Upon further Ag+ addition, after half of an equivalent Ag+ per equivalent CN-, the precipitation of
AgCN(s) occurs.

 2 AgCN( s )
[ Ag (CN )2 ]  Ag  

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):
32
0
E  EAg


/ Ag
RT
ln[ Ag  ]
nF
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. This is based on
the relation between equilibrium constant and electrode potential:
nEF  RT ln K
Preparation:











0.2 M AgNO3solution
1 M K2CrO4
KSCN
Fe(NO3)3 9H2O
KCl
KBr
HNO3 conc.
0.1 % Eosin solution
Burette 50 ml
Magnetic stirrer and stirring rod
Wide-necked conical flask
Procedure:
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) nitrate, Fe(NO3)3 9H2O, are transferred
into a large test tube and diluted with a few ml of water. Some crystals of ammonium thiocyanate
33
NH4SCN are dissolved in a medium test tube in a few ml of water. This solution is added drop wise 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 drop wise. The indicator is a 1 M
K2CrO4. What causes the brownish colour? Why is this compound formed only when all Br- has 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 eosin 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
The amount of solution made according to description here is sufficient for 3-4 students. Weigh
about 4.4 g of AgNO3 into a large test tube, dissolve with some water and pour the solution into a
500 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.
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.
34
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-, SCN-
Carry 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
2
1
5
7
3
11
4
8
3
9
2
6
7
8
9
1
10
Titration setup
35
Fractioning 2
Condensation, distillation, sublimation
The fractioning methods discussed in this chapter are based on the different volatilities of the
substances to be separated. Volatility is determined by the dependence of vapour pressure as a
function of temperature, which also defines the boiling point. Such processes are important in
nature, especially concerning the “elixir of life”, namely water. Realms of different physical states of
the same compound being in contact are called “phases” by scientists. On Earth’s surface we are
always within the phases of liquid and gaseous water. We contain liquid water ourselves; the air
around us contains water vapour, and in the lung we have water vapour at body temperature. If the
air outside is considerably colder than the body, condensation sets in upon exhalation, and we can
see a kind of fog, an aerosol. If the air is warm and also saturated with water vapour, water released
by our skin cannot evaporate: we sweat. If the air becomes drier, the water will evaporate, and the
skin is cooled by this. Therefore, we feel more comfortable at 40°C in the desert than at 30°C on a
damp day at or latitude, where there is more water in the environment. In presence of the solid
phase, ice, the vapour pressure of water is very low. Air in contact with ice and snow is as dry as in a
desert. In regions extremely deficient in water, where no rain may fall for years (North Chile,
Namibia) condensation processes (fog and dew) allow the existence of frugal plants and animals.
Rain and snow are both caused by condensation of water vapour by cooling in the atmosphere.
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. This principle is also called the “cold
trap”. In a distillation the evaporated substances must be re-condensed to liquid, which is rather
difficult on the laboratory scale with very volatile materials.
36
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.
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.
37
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
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.
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 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
38
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.
Distillation of an 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.
Preparation:
-
Set up the distillation apparatus based on ground glass joint according to the above drawing.
A water bath (400-600 ml beaker) heated by gas burner or electric plate serves as the heat
source. Attention: ethanol is flammable!
-
About 50 ml of wine are transferred into the distillation flask.
-
Boiling aid is added, the cooler is attached and the coolant flow turned on.
-
Pasteur pipette with balloon and felt tip pen are kept ready.
39
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
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 drawn into a Pasteur pipette such that no air is taken up from
below. The liquid level in the pipette is marked with the felt tip pen. 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 drawing water into the pipette to the mark and weighing this amount of
water like before (assumption: ς(H2O) = 1 g ml-1). Continue the distillation and note the temperature
changes with time.
40
Density
%weight %vol
Density
%weight %vol
Density
%weight %vol
20°C, g/ml Ethanol Ethanol 20°C, g/ml Ethanol Ethanol 20°C, g/ml Ethanol Ethanol
1
0
0
0.94662
35
41.9
0.87158
69
76
0.99813
1
1.3
0.94473
36
43
0.8692
70
76.9
0.99629
2
2.5
0.94281
37
44.1
0.8668
71
77.8
0.99451
3
3.8
0.94086
38
45.2
0.8644
72
78.6
0.99279
4
5
0.93886
39
46.3
0.862
73
79.5
0.99113
5
6.2
0.93684
40
47.4
0.85958
74
80.4
0.98955
6
7.5
0.93479
41
48.43
0.85716
75
81.2
0.98802
7
8.7
0.93272
42
49.51
0.85473
76
82.1
0.98653
8
10
0.93062
43
50.6
0.8523
77
83
0.98505
9
11.2
0.92849
44
51.6
0.84985
78
83.8
0.98361
10
12.4
0.92636
45
52.6
0.8474
79
84.6
0.98221
11
13.6
0.92421
46
53.7
0.84494
80
85.4
0.98084
12
14.8
0.92204
47
54.7
0.84245
81
86.2
0.97948
13
16.1
0.91986
48
55.8
0.83997
82
87.1
0.97816
14
17.3
0.91766
49
56.8
0.83747
83
87.9
0.97687
15
18.5
0.91546
50
57.8
0.83496
84
88.7
0.9756
16
19.7
0.91322
51
58.8
0.83242
85
89.5
0.97431
17
20.9
0.91097
52
59.8
0.82987
86
90.2
0.97301
18
22.1
0.90872
53
60.8
0.82729
87
91
0.97169
19
23.3
0.90645
54
61.8
0.82469
88
91.8
0.97036
20
24.5
0.90418
55
62.8
0.82207
89
92.5
0.96901
21
25.7
0.90191
56
63.8
0.81942
90
93.2
0.96763
22
26.9
0.89962
57
64.8
0.81674
91
94
0.96624
23
28.1
0.89733
58
65.8
0.81401
92
94.7
0.96483
24
29.2
0.89502
59
66.8
0.81127
93
95.4
0.96339
25
30.4
0.89271
60
67.7
0.80848
94
96.1
0.9619
26
31.6
0.8904
61
68.6
0.80567
95
96.7
0.96037
27
32.7
0.88807
62
69.6
0.8028
96
97.4
0.9588
28
33.9
0.88574
63
70.5
0.79988
97
98.1
0.95717
29
35.1
0.88339
64
71.5
0.79688
98
98.7
0.95551
30
36.2
0.88104
65
72.4
0.79383
99
99.3
0.95381
31
37.4
0.87869
66
73.3
0.79074
100
100
0.95207
32
38.5
0.87632
67
74.2
0.95028
33
39.6
0.87396
68
75.1
0.94847
34
40.7
41
Volatile Compounds
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.
The most important volatiles in biology are water, oxygen, nitrogen and carbon dioxide. Further
examples are volatile organics which are secreted by creatures. These comprise pheromones and
metabolism products like ethanol or methane. Further “biogases” are ammonia, hydrogen sulphide
and laughing gas (dinitrogen monoxide). Water is the essential solvent in all biochemical reactions.
There exist hydrophobic compartments in creatures, however, they are not strict anhydrous in a
physical-chemical sense.
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
42
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.
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.
43
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.
44
Determination of the molar mass by melting point depression (cryoscopy)
Introduction
A dissolved compound decreases the melting point of a solvent compared to its pure state. The
extent of this melting point depression is determined by the reciprocal heat of melting 1/ΔHmelt on
the side of the solvent. Solvents with small heats of melting show the strongest effect. On the side of
the solute it is the number of dissolved molecules per number of solvent molecules that determines
the extent. This ratio is proportional to the molality, a concentration measure with symbol b.
For dilute solutions we can write:
ΔTf = λf •b•z
ΔTf : melting point depression
λf : cryoscopic constant, molar melting point depression
b : molality of the solute = moles of solute per kg of solvent
z : number of particles formed in solution per molecule of solute, with salts > 1
If we dissolve the mass mg of a compound in the mass mL of a solvent, we have:
mg
ng = M
g
ng
mg
Tf = f m z = f m M z
L
L g
with M: relative molar masses, n: number of moles, mg mass of solute in g , mL: mass of solvent in kg.
We obtain: Mg = f
mg z
mL Tf
Reagents:

Assemble apparatus as depicted, perforate cap with scissors

Electronic thermometer: test function by immersion into ice/water mixture, this
should show 0 °C. Replace battery if necessary
45
Procedure
The sensor tip of an electronic thermometer (0.1 °C resolution) is inserted
into a pill tube with perforated plastic cap. The temperature difference to
the environment is kept small by immersing the “instrument” repeatedly
into the cooling bath for a moment.
In order to determine the molar mass of acetone we use the ether solvent
1,4-dioxan which melts at 11.74 °C has a λf = 4.63 K kg mol–1. The dry and
weighed apparatus is filled with 1,4-dioxan to 3-4 cm height, weighed
again, and the melting point is determined. For the cooling, an ice/water
mixture is sufficient. The pill tube is immersed and slightly shaken, until a
few solvent crystals form. Never let the solvent freeze thoroughly, it
would be a waste of time! The tube is removed from the bath and the
melting of the crystals is observed. If the melting is too rapid, the tube is
cooled again for a short time. We read the temperature when the last
remaining crystal is just dissolving. We cool again and test the reproducibility of the process. In case
of doubt, repeat several times. Never let the liquid come to rest, because this would lead quickly to
inhomogeneous temperature distribution.
Add about 0.5 ml of acetone to the 1,4-dioxan and weigh again. Determine the melting point of the
mixture as before with the pure solvent. From the change of melting point and the weights,
determine the molar mass of acetone, as described above.
46
Acids and Bases
Acid and base definitions by Lewis and Brønsted-Lowry
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

 BF4
BF3  F  

(Lewis)

 NH 4  Cl 
HCl  NH3 

(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
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.
Acid-base equilibria are essential in biology. Organisms are adapted to the acidity of their
environment; large creatures actively regulate their internal acid-base balance. The general degree of
acidity on Earth’s surface is mainly determined by water, which itself has acid and base properties.
47
This is different from planets like Venus, where the dominating sulphuric acid creates an entirely
different ambience, or the gas planets, where ammonia is abundant. The degree of acidity of water,
which we characterise by the pH value, is modulated by solutes. Most significant is the combination
of carbon dioxide and carbonates, which governs the pH of the sea and the internal pH of most
creatures on land. The sea can take up and release carbon dioxide from or to the air. As counterpart,
carbonates are formed or dissolved. Organisms degrade organic material by cellular respiration to
carbon dioxide and water. Carbon dioxide is eliminated by the respiratory organs. An organism can
control its internal pH by regulating its carbon dioxide expiration rate. The stabilisation of internal pH
is important because most biocatalysts, called enzymes, have limited pH ranges of operation. This is
not necessarily neutral; there exist microorganisms which live in acidic water in volcanic areas. On
the basic side life cannot go that far, because bases catalyse the decomposition of proteins by water
much better than acids.
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 …
48
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:

 H 3O  B 
HB  H 2O 

(1)
The law of mass action for this reaction can be written as:
[ H 3O  ][ B  ]
K'
[ HB][ H 2O]
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.
In dilute aqueous solutions, [H2O] = constant, because it is usually in high excess.
With K' [H2O] = K the law of mass action can be simplified to:
49
K
[ H 3O  ][ B  ]
[ HB]
(2)
With the definitions pK = -log K and pH = -log[H3O+] the result, in logarithmic form is:
pH  pK  lg
[ B ]
[ HB]
(3)
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

 H 3O  OH 
H 2O  H 2O 

(4)
50
transfers a proton to itself. The law of mass action applied to this reaction yields
[ H3O ][OH  ]  K w
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
Concept of pH
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 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
pH  pK  lg
[ B]
[ HB]
51
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).
pH  pK  lg
[ Ind ]
[ HInd ]
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 having different pK values, so-called universal indicators, we can
quickly estimate pH of solutions over the entire pH range.
Synthesis of ethanolic hydrogen chloride
H2SO4: proton transfer to Cl-. Formation of HCl(g).
52
In this 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.
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.
vacuum adapter
PVC tubing
Bent joint
Ethanol
Dropping
funnel
Magnetic
stirrer
5
4
H2SO4
6
2
1
5
7
3
4
8
3
9
2
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+ + Cl-
O
Ethanol
+ NaCl(s)
OH
53
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
solution is poured into a crystallisation dish and the ethanol is allowed to evaporate (in a ventilated
hood).
Esterification of boric acid
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
H 2 SO4

 B(OH )3
CH3OH

B(OCH3 )3
Reaction between gaseous NH3 and HCl
Ammonia is usually sold 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:

 NH3  H 2O
NH 4  OH  

Liquid ammonia is a water-like solvent, though less acidic and with some interesting properties.
NH3 is a volatile weak base. In a fume hood, Hold the open bottles of concentrated ammonia and
concentrated hydrochloric acid solutions close to each other for a short time and observe.
54
Synthesis of a calcium salt
CaCO3 is a very convenient starting material for the preparation of other calcium compounds from
free acids. It is available in high purity degrees, stoichiometric 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
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.
Sublimation of ammonium chloride
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 gas burner.
Aluminium chloride as Lewis acid
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
55
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.
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.
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 non-shining
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 it 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
56
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.
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 + SO42CO2 + H2O + CO32-
2 HSO42 HCO3-
and also the formation of hydrogen tartrate from tartaric acid and its sodium potassium mixed salt.
Both are highly water soluble while the monopotassium salt is sparingly soluble. Tartatric acid occurs
to have 3 structural isomers with different symmetry. The subject is more profoundly taught in the
organic chemistry classes.
Weigh 4.0 g tartaric acid and 7.9 g sodium potassium 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 re-obtain the mixed cations salt
KNa(tart)4 H2O, also called Seignette's salt.
57
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 expand their
coordination numbers under addition of the oxygen in water release protons thereby, at a wide
range of acidity.

SO3  H 2O 
 H 2 SO4 
 H aq
 HSO4aq
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:

OH
SO2 ( aq )  H 2O 
 HSO3  H  
 H 2O
In the same way water acts as a base against non-metallic halides and oxohalides:
PCl3  2H 2O 
 H3 PO3  3HCl
COCl2  H 2O 
 CO2  2HCl
Highly charged metal aquo ions, as they occur normally in aqueous solution, can liberate protons
from the water they bind:

 [ Al ( H 2O)5 (OH )]2  H 
[ Al ( H 2O)6 ]3 

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:
58
O2  H 2O 
 2OH 

 OH   HCO3
CO32  H 2O 


 OH   HF
F   H 2O 

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
Write down the reactions with water which cause the observations.
Acidimetric titration
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:

 2 H 2O
H3O  OH  


 B   H 2O
HB  OH  

If we plot, as a result from such a titration, 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,
59
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.
Molar mass M of the acid 
m(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.
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.
* Use 0.1 M NaOH calibrated solution which is commercially available ("Titrisol").
** Later, compare the neutralisation curves recorded and determine the suitable pH indicator for the
corresponding titration.
Determination of the pK values of the indicator thymol blue
Thymol blue is an indicator which can release two protons stepwise:
pK1
pK2



 H   Ind 2
H 2 Ind 
H   HInd  

60
red
yellow
blue
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).
For the estimation of pK2, the following solutions are prepared in 6 large test tubes:
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
Estimate the pH value for test tubes 1-5. Remark: in tubes 2-3 mixtures of NH4+/NH3 with
concentration ratios of approximately 10:1, 1:1 and 1:10 are formed. The pK of NH4+ can be found in
tables (e.g. in the appendix here). Then
pH  pK  lg
[ NH 3 ]
[ NH 4 ]
Add indicator solution and estimate pK2 of thymol blue. What is the approximate pH of tap water?
61
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.
I
1
ci Zi2

2 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.
62
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.
Calculation: from pH  pK  lg
[ B]
1
2
the ratio [B]/[HB] can be obtained. With I   ci Z i we
[ HB]
2 i
can determine [B] and [HB] for the required ratio. Finally, with [B] and [HB] known, we can calculate
the needed amount of both phosphates. Measure the pH of the buffer prepared wit a glass electrode
and the pH meter.
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.
63
Glass Electrode
Combined
Glass Electrode
Connecting reference
electrode:
Usually Ag/AgCl or Pt
3 M KCl
Glass membrane:
Exchanges
+
+
+
Na /K and H
Porcelain
frit
[H+ ] inside
Buffer pH = 7
in 3 M KCl
+
[H ] inside
+
[H ] outside
+
+
+ +
+
+ + +
++
++
+ + +++ + +
+ +
+
+ +++ + +
+
+
+
External reference
electrode:
Ususally Ag/AgCl
+=H
+
The ratio of occupation of in- and
outside of the glass membrane
by H+ generates an electrostatic
potential. Anions are not bound to
the exchanging groups at the silicate
scaffold. Occupation on the inside is
kept constant by a buffer solution
such that only the occupation on the
outside determines the potential.
Other solid membrane electrodes
work by the same principle.
+
In a combined glass electrode a
second electrochemical half-cell
is included. It is called the reference
electrode. Usually it is integrated
in an envelope of the tube that
bears the glass membrane sphere.
It contains a Ag/AgCl reference
electrode in 3M KCl. The contact to
the solution under examination is
established by a porous frit just
above the glass sphere.
+
E = E°' + RT/F * ln ([H ] aussen/[H ] innen)
64
Redox Reactions
Redox reactions, solvent free or in aqueous solutions
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.
In biology redox reactions are important for conversion and release of energy. Cellular respiration
and photosynthesis are coupled redox systems at the basic level, before energy is distributed in the
form of ATP (adenosine triphosphate), which undergoes hydrolysis. Redox reactions are catalysed by
many other enzymes not involved to cellular respiration. Catalase decomposes hydrogen peroxide,
peroxidases use it to oxidise small molecules, monooxygenases hydroxylate organic molecules by
using dioxygen. The purpose is often to make these more water-soluble in order to excrete them. An
example for the conversion of a toxic compound is the oxidation of ethanol by alcohol
dehydrogenase. However, the first intermediate, acetaldehyde, is even more toxic than ethanol.
Fortunately it is quickly oxidised further to acetic acid which is not toxic.
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 )6 ]4 
 Ce3  [ Fe(CN )6 ]3
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:
65
MnO4  8H   5e 
 Mn2  4H 2O
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
reaction type with the synthesis of KICl4, the reaction of IO3- with I- and the permanganometric and
iodometric titrations in the following experiments.
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 )O3(  II )  
 Cl (  I )  O2(0)
In order to keep stoichiometry 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
66
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
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-.
67
Preparation of CuCl
The chloride of monovalent copper CuCl can be obtained by the reduction of Cu2+ with elemental Cu.
Formally this is the reaction
Cu 2  Cu 0 
 2Cu 
followed by the precipitation
Cu   Cl  
 CuCl
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 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, keep on vacuum until dry.
Redox reactions in qualitative analysis
Redox system Fe(II) / Fe(III) - Sn(II) / Sn(IV)
Fe3+ can be reduced to Fe2+ by Sn(II) under formation of Sn(IV):
2Fe3  Sn II 
 Fe2  Sn IV
68
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).
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.
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).
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.
Mix 5 parts of potassium iodide with 1 part of potassium iodate (moles), such that the total mass is
about 0.2 g and dissolve in water. It is important to dissolve the crystals completely (why?) Add a few
drops of 2 M HCl.
69
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 potato are cut into small pieces and these are 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 stirring with a glass rod. (Never use metal, e.g. a spatula, for this operation). The extract is
decanted and centrifuged.
Transfer 2 ml of potato extract into each of two test tubes, and 2 ml Cu2+ solution into each of two
others. A drop Na2S solution is added to one of the tubes with extract and to one with Cu2+. 3 ml of
freshly prepared 1 M H2O2 are added to all test tubes. Observe. What is happening here?
Standard reduction potential Fe(CN)63- / Fe(CN)64Another 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.
70
Examples:
2H   2e (Pt) 
 H2
reduction half cell at inert cathode
Zn0 
 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+.
Standard 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)64- in
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.
71
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 )6 ]4
[ Fe(CN )6 ]3  e 

E  Eref  E
0
[ Fe ( CN )6 ]3 /[ Fe ( CN )6 ]4
3
RT [ Fe(CN )6 ] 

ln
F
[ Fe(CN )6 ]4 
Enter the corresponding values of E, [Fe(CN)6]3-, [Fe(CN)6]4- into the equation and calculate
E0([Fe(CN)6]3-/Fe(CN)64-).
Permanganometric titration
The 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".
72
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.
Procedure
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.
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  
 2Fe2  I 2
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:
73
I3  2S2O32 
 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
2Cu 2  4I  
 2CuI ( s )  I 2
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.
74
Ligand Exchange and Complex Formation
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  5H 2O  4 NH3 
 [Cu( NH3 )4 ]SO4  H 2O  4H 2O
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 stoichiometric 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
tetrahedral
CuCl42- tetrachlorido cuprate
Cu
+II
Cl
4
quadratic
AlF63- hexafluorido aluminate
Al
+III
F
6
oktahedral
Ni(CO)4 nickel tetracarbonyl
Ni
0
C
4
tetrahedral
SbF5 antimony pentafluoride
Sb
+V
F
5
trig. bipyr
Ag(S2O3)23- bis-(thiosulfato) silver
Ag
+I
S
2
linear
75
Illustration of the most important coordination geometries: the bonding donor atoms are located at
the corners of the polyhedron, the metals represented by the sphere at the centre.
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 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 often 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.
76
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 ( H 2O) x ]n  yL 
 [ MLy ( H 2O) x  y ]( n y )  yH 2O
while the replacement of a ligand L by another ligand is called "ligand substitution"
MLx  yB 
 MLx  y By  yL
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 variable the interplay of all facts influencing complex formation is, as diverse is 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 that form sparingly soluble precipitates with metal ions are useful
in gravimetry. Water softeners bind calcium. Complexones are used for the titration of various metal
ions. The red blood colour is an inert iron complex for the transport of O2. Vitamin B12 is a cobalt
complex. Complex formations are used as detection methods in qualitative analytical chemistry.
Complex catalysts promote polymer synthesis. Chlorophyll is a magnesium complex, etc.
Biologically important are complexes of Mg2+, Ca2+, Mn(2-4)+, Fe(2-4)+, Co+, Cu(1-2)+, Zn2+, Mo(4-6)+. Some
microorganisms contain enzymes based on vanadium and nickel. Biological ligands are often derived
77
from amino acids. Proteins and peptides contain amino acids that provide donor functions, e.g.
carboxylate from aspartate and glutamate, phenolate from tyrosine, imidazol from histidine and
thiolate from cysteine. The oxygen-containing donors bind Mg2+, Ca2+, Mn(2-4)+ and Fe3+. The others
bind most metals except Mg2+, Ca2+ and Mn(2-3)+. Inorganic helper ligands are hydroxide and oxide
which often form bridges between several metal atoms. This kind of complex is called a cluster.
Another bridging ligand is sulphide. It forms functionally important aggregates with Fe(2-3)+. A special
class are the macrocyclic ligands. They confer unusual properties to metals regarding their Lewis acid
and redox properties. The largest group in biology are the porphyrins, and thereof the Fe complexes
which are called haemes. They are related to the chlorophylls which are crucial in photosynthesis,
and to corrins which form Vitamin B12 with Co+. Vitamin B12, cobalamin, is essential for the transfer of
methyl (CH3) groups.
O
H2N
O
H2N
CH3
H3C
O
H2N
NH2
N
H3C
O
N
H3C
N
+
Co
N
N
CH3
O
NH2 H C
3
H3C
O
O
NH2
O
O
NH
CH3
CH3
N
CH3
P
O
O
-
HO N
CH3
H
H
H
HO
O
H
Vitamin B12 (Cobalamin)
Haem complexes carry out the functions of oxygen transport (haemoglobin), oxygen storage
(myoglobin) and many redox reactions (peroxidases, oxygenases). Ion selective ligands that bind
selectively monovalent cations can also be found in biology. They enable controlled transport of ions
across lipophilic membranes. Valinomycin is a representative:
78
Valinomycin
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. CrO2Cl- and
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 CrO3Cl-
CrO42  Cl   2H  
 CrO3Cl   H 2O
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  
 CrO3Cl2  2H 2O
If CrO3Cl- or CrO2Cl2 are brought into water the reactions are reverted.
79
1 g of potassium dichromate K2Cr2O7 is 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
80
A metal ions (Al3+, Ti4+) prefer OH- as a ligand over NH3: addition of NH3 causes metal hydroxide
precipitation.
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:
Ni 2  xNH3 
 [ Ni( NH3 ) x ]2
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 )  4 NH3 
 [Cd ( NH3 )4 ]2  2OH 
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.
81
Al (OH )3 ( s )  3OH  
 [ Al (OH )4 ]
Zn(OH )2 ( s )  2OH  
 [ Zn(OH )4 ]2
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.
e) Inertness of Fe(CN)64-
The complex particle Fe(CN)64- should decay into Fe2+ and hydrocyanic acid upon acidification:
[ Fe(CN )6 ]4  4H  
 Fe2  6HCN
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 )6 ]4  4H  
 H 4 [ Fe(CN )6 ]
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.
f) 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+,
82
-O 3S
OH
-O 3S
OH
can bind totally three of these tiron anions L2-. In the formation process, according to
Fe3  H 2 L 
 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 FeL25- dominates
and at pH  10 the red 1:3 complex FeL39- prevails.
1 mg of Fe(NO3)39H2O is weighed into a 50 ml beaker, dissolved in 30 ml water and acidified with
one drop of 2 M HCl. Add 10 mg of tiron. The blue 1:1 complex appears. Take a sample of this
solution in a polystyrene cuvette. By addition of solid CH3COONa3H2O to the solution in the beaker,
the pH is increased until the colour changes to violet (1:2 complex). Take a sample of this solution in
a second cuvette. Upon addition of a drop of concentrated NH3 to the beaker the red 1:3 complex is
formed. Take another sample in a third cuvette. 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).
Fill a forth cuvette with water and take the reference spectrum from 350 nm to 700 nm. Measure the
according spectra of all three iron-tiron complex samples. Compare the wavelength of the absorption
maxima of the samples with their colour of appearance. Is there any relationship?
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
83
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, cis-Co(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-.
84
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.
85
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 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.
86
Dye
pH/(reagent)
Metal
Murexide
14 (NaOH)
Cd2+
Erio T
9-10 (NH3/NH4+)
Ca2+
Pyridyl-azo-naphthol (PAN) 5 (CH COOH/CH COO-) Cu2+
3
3
Note the colours before and after metal addition in a table. Finally, add 0.1 M Komplexon III solution
drop wise 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.
Determination of the hardness of water by complexometric titration
For a polydentate ligand which saturates multiple or even all of the coordination sites of a metal ion
a very simple stoichiometry 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 stoichiometry
M  Hi L 
 ML  iH 
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
87
H4EDTA, abbreviated H4Y
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 Eriochrome black T ground with NaCl 1:100.
Why is the buffer required?
88
50 ml of 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 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
89
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 the silica gel and the
individual dyes are transported upwards at different velocities. When the solvent front has reached
90
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
91
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
2Br  
 Br2  2e
oxidation
CH3COO  H  
 CH3COOH protonation
Al 3  3HOxin 
 Al (Oxin)3  3H 
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 chlorido 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.
92
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.
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
K I2 
with
[ I 2 ]H 2 O
[ I 2 ]CH 2Cl2
[ I 2 ]H 2O 
n( I 2 ) H2O  0.2 l  [ I 2 ]H2O
0.01M Vtitr
2 150 ml
and
[ I 2 ]CH 2Cl2 
n( I 2 )tot  n( I 2 ) H 2O
0.01 l
n( I 2 )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.
93
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:
weaker
Li+< H+ < Na+ < Cs+ < Mg2+< Ca2+< Al3+< Ce3+
stronger
weaker
F-< Cl-< Br-< NO3- < HSO4- < I-
stronger
If an ion exchanging resin RA+ were transferred into a solution of ion B+ and thoroughly mixed, the
equilibrium

 RB  A
RA  B 

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
94
complexes on an anion exchanger. The separation does not succeed with commercially available
cation exchangers.
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.
95
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 9 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 9M HCl and pour also into the tube. Knock gently at the tube, 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 9 M HCl to ensure that it is in the Cl- form. Close the
stopcock when the NaCl 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 9 M HCl and drain the column until the
upper level is again 1 cm above the glass wool ball. Add 15 ml of 9 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 4.5 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 1 M or 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+.
96
Qualitative analysis
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 yellowbrownish 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 and is also
97
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 we start 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.
Samples are provided by the teaching assistants. They may contain all or only part of the ions
discussed above.
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.
98
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.
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
99
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 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
100
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 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!
101
-
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 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.
102
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
103
Appendix
pk values of some acids at 25°C
Säure
pKs-Werte
Ammonium (NH4+)
9.23
Boric acid (B(OH)3)
9.14
Acetic acid (HOAc, CH3COOH)
4.75
Ethylenediamine tetraacetic acid (H4EDTA, H4Y)
2.0
2.69
Glycinium (Hgly+, NH3+CH2COOH)
2.35
9.78
Carbonic acid (H2CO3)
6.37
10.25
Oxalic acid (H2ox, HOOCCOOH)
1.23
4.19
Phosphoric acid (H3PO4)
2.12
7.21
Nitric acid (HNO3)
-1.43
Hydrochloric acid (HCl)
-6.1
Sulphuric acid (H2SO4)
-8.0
1.92
Dihydrogen sulfide(H2S)
7.04
19
„Htris+“ (NH3+C(CH2OH)3)
8.09
6.18
10.15
12.67
104
pH indicators
Indicator
Acidic colour
pH range of colour
change
pK
Alkaline colour
Alizarine yellow
yellow
10.4
-
12.0
11
pink
Bromocresol green
yellow
3.8
-
5.4
4
blue
Bromocresol
yellow
5.2
-
6.8
5
violet
Bromothymol blue
yellow
6.0
-
7.6
7
blue
Cresol red
yellow
0.4
-
1.8
yellow
7.0
-
8.8
8
violet
Litmus
red
4.4
-
6.2
6
blue
Methyl red
red
4.8
-
6.0
5
yellow
Neutral red
red
6.8
-
8.0
7
yellow
colourless
8.2
-
10.0
9
red
Phenol red
yellow
6.6
-
8.0
7
violet
Thymol blue
red
1.2
-
2.8
2
yellow
yellow
8.0
-
9.6
9
blue
purple
Phenolphthalein
red
Further the following indicators are used:
for complexometry:
acidic colour
basic colour
Murexide
violet
blue
Erio T
violet-brown
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.
105
Standard reduction potentials
E0 (V)
Ag+/Ag .................................................................. 0.7996
AgCl/Ag .................................................................. 0.2223
Br /2 Br- ............................................................... 1.065
2
4+
Ce /Ce3+ (1 M H2SO4) ................................... 1.4587
Cl /2 Cl- ................................................................ 1.3583
2
3+
Co /Co2+ (3 M HNO3)..................................... 1.842
[Co(NH3)6]3+/[Co(NH3)6]2+................................ 0.1
Cr3+/Cr2+ ............................................................. -0.41
Cr O 2-/2 Cr3+ .................................................... 1.33
2 7
+
Cs /Cs .................................................................. -2.923
Cu+/Cu .................................................................. 0.522
Cu2+/Cu+ .............................................................. 0.158
Cu2+/Cu ................................................................ 0.3402
Fe3+/Fe2+.............................................................. 0.770
[Fe(CN) ]3-/Fe(CN) ]4- (1 M H SO ) ............. 0.69
6
6
2
4
+
2 H /H2.................................................................. 0.0000
H2O2/2 H2O ........................................................ 1.776
Hg2Cl2/2 Hg (Kalomel) (satd. KCl) ............... 0.2415
I /2 I- ..................................................................... 0.535
2
+
K /K ...................................................................... -2.924
Li+/Li ..................................................................... -3.045
MnO2/Mn2+ .......................................................... 1.208
MnO -/MnO 2- .................................................... 0.564
4
4
MnO4-/MnO2 ....................................................... 1.679
MnO -/Mn2+ ......................................................... 1.496
4
+
Na /Na ................................................................. -2.7109
Ni2+/Ni.................................................................. -0.23
O2/2 H2O ............................................................. 1.229
Rb+/Rb ................................................................. -2.925
S/S2- ...................................................................... -0.508
S4O62-/2 S2O32- ................................................. 0.09
Zn2+/Zn ................................................................ -0.7628
106
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
Mg2+
20
0.1 M KNO3
5.6
Ca2+
25
‘’
4.72
20
‘’
4.2
25
‘’
3.8
2+
Sr
Ba
2+
Al3+
20
‘’
14.8
3+
22
0.5 M NaCl
21.84
3+
20
0.1 M KNO3
13.0
Mn2+
20
‘’
11.7
2+
20
‘’
10.7
Co2+
20
‘’
13.8
Sc
La
Fe
2+
20
‘’
17.4
2+
20
‘’
19.6
Zn2+
20
‘’
13.1
2+
20
‘’
10.4
2+
Hg
25
‘’
17.50
Pb2+
20
‘’
13.5
Ni
Cu
Cd
b) NH3 complexes (pK of NH4+: 9.24, 25°C, I
M
T (°C)
0)
Medium
log K1
log K2 log K3 log K4 log K5 log K6
2 M NH4NO3
2.11
1.63
1.05
0.76
0.18
-0.62
0.65
0.08
Co2+
30
Ni2+
‘’
‘’
2.78
2.27
1.65
1.31
Cu2+
‘’
‘’
4.14
3.52
2.87
2.15
Zn2+
‘’
‘’
2.45
2.28
2.64
2.11
Ag+
25
‘’
3.35
3.90
107
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
Cu2+
25
Ni2+
‘’
0.1 M KNO3
‘’
log K1 log K2 log K3
8.23
6.96
5.73
4.83
3.44
108
Solubility products
MaLb(s)
a M+ b L
KL = [M]a.[L]b
Solid
pKL (at 25°C)
AgCl
9.80
AgBr
12.27
AgI
16.07
Ag2S
50.22
CdS
25.10
CuS
36.22
FeS
18.22
HgS
52.70
MnS
13.52
NiS
20.97
PbS
27.52
ZnS
24.70
109
Conductivity data
Molar asymptotic conductivity 0 of ions in water at 25°C (in Scm2mol-1)
H+ .............. 349.65
Ag+ .................61.9
La3+............. 209.1
Na+ .............. 50.08
Ba2+ ............ 127.2
N(C2H5)4+ ........... 32.6
K+................. 73.47
Co2+ ............110
Co(NH3)63+ . 305.7
OH- ................. 198
NO3- .............71.42
SO42- ............... 160
Cl- ................ 76.31
C2O42- ....148.22
ClO4- ................... 67.3
Br- .................. 78.1
CO32- ........138.6
Fe(CN)63- .... 302.7
Temperature dependence of 0: example HCI
T (°C)

0 (Scm2mol-1)
5
15
25
35
45
55
296.4
360.8
424.5
487.0
547.9
606.6
Concentration dependence of : example NaCI
c (moll-1)

 (Scm-2 mol-1)
0
5.10-4
10-3
5.10-3
10-2
10-1
126.39
124.44
123.68
120.59
118.45
106.69
Some values of 0 in non-aqueous Media at 25°C (Scm2 mol-1)
Solvent
NaCI
KCI
KBr
CH3OH
98
105
109
C2H5OH
42
45
Dielectric constant
32.6
24.3
110