laboratory experiments qualitative analysis

Department Of Biochemistry and Medical Chemistry

Medical School, University of Pécs

Qualitative chemical analysis

from

LABORATORY EXPERIMENTSLABORATORY EXPERIMENTSLABORATORY EXPERIMENTSLABORATORY EXPERIMENTS

IN MEDICAL CHEMISTRYIN MEDICAL CHEMISTRYIN MEDICAL CHEMISTRYIN MEDICAL CHEMISTRY

Edited by GYÖRGY OSZBACH

PÉCS, 1995

CONTENTS

1. INTRODUCTION 3

Laboratory safety..................................................................................................... 3

Accident protection, fire protection, first aid .......................................................... 5

2. LABORATORY UTENSILS AND METHODS 9

Utensils.................................................................................................................... 9

Simple laboratory methods...................................................................................... 9

Laboratory measurements .................................................................................. 9

Heating ............................................................................................................... 9

Forming precipitates ........................................................................................ 12

Preparation of gases......................................................................................... 12

3. SIMPLE LABORATORY SEPARATION METHODS 14

Decantation............................................................................................................ 14

Filtration ................................................................................................................ 14

Drying.................................................................................................................... 15

4. CHEMICAL ANALYSIS 17

4.1 QUALITATIVE CHEMICAL ANALYSIS .......................................................... 18

Detection of cations............................................................................................... 19

Detection of anions................................................................................................ 32

Simple analysis of cations and anions ................................................................... 39

Self-test questions ................................................................................................. 42

9. APPENDIX 44

9.1 INORGANIC CHEMISTRY; A short overview................................................... 44

Classification of the inorganic compounds ........................................................... 45

The nomenclature of the inorganic compounds .................................................... 51

Solubility of inorganic compounds in water ......................................................... 55

9.3 CONCENTRATIONS OF REAGENT SOLUTIONS .......................................... 56

INDEX OF LABORATORY EXPERIMENTS

2. LABORATORY UTENSILS AND METHODS

2.1 Preparation of a solution of a given concentration.................................................12

2.2 Preparation of capillary tubes.................................................................................15

3. SIMPLE LABORATORY SEPARATION METHODS

4. CHEMICAL ANALYSIS

4.1 QUALITATIVE CHEMICAL ANALYSIS

Detection of cations

4.1.1-5 Reactions of lead(II) ion.............................................................................34

4.1.6-9 Reactions of mercury(I) ion........................................................................35

4.1.10-13 Reactions of mercury(II) ion ......................................................................36

4.1.14-18 Reactions of copper(II) ion.........................................................................37

4.1.19-21 Reactions of arsenic(III) ion .......................................................................38

4.1.22-25 Reactions of cobalt(II) ion..........................................................................40

4.1.26-28 Reactions of chromium(III) ion..................................................................41

4.1.29-33 Reactions of zinc ion..................................................................................42

4.1.34-36 Reactions of iron(II) ion .............................................................................43

4.1.37-40 Reactions of iron(III) ion............................................................................44

4.1.41-43 Reactions of calcium ion ............................................................................45

4.1.44-46 Reactions of magnesium ion ......................................................................46

4.1.47-48 Reactions of sodium ion.............................................................................47

4.1.49-51 Reactions of potassium ion ........................................................................47

4.1.52-54 Reactions of ammonium ion ......................................................................48

4.1.55-56 Reactions of hydrogen ion..........................................................................49

Detection of anions

4.1.57-59 Reactions of carbonate ion .........................................................................51

4.1.60-61 Reactions of hydrogen carbonate ion .........................................................52

4.1.62 Reaction of hypochlorite ion ......................................................................52

4.1.63 Reaction of sulphate ion.............................................................................53

4.1.64-66 Reactions of phosphate ion ........................................................................53

4.1.67 Reactions of chloride ion............................................................................54

4.1.68-70 Reactions of iodide ion...............................................................................55

4.1.71-72 Reactions of nitrite ion ...............................................................................56

4.1.73-74 Reactions of nitrate ion ..............................................................................57

4.1.75 Reactions of hydroxide ion ........................................................................57

4.1.76 Simple analysis of cations and anions........................................................58

PREFACE

ince the English Program started at the University Medical School of Pécs in

1984, remarkable changes have been made in chemical education. Today,

there are appropriate textbooks available to support the lectures and seminars

in the fields of General, Organic and Bioinorganic Chemistry. In 1994, the

Department issued a handout on Inorganic Chemistry, but there was a lack of an

updated laboratory manual. (The former one was written in 1984 and the last reprint

was edited in 1991.) This edition is a translation of the Hungarian version ("Orvosi

kémiai gyakorlatok") issued this year.

The manual contains experiments much more than can be completed within

the limited time of the main course. Among them, there are practices to be replaced in

the future by another ones that cannot be performed for the time being. Some of the

experiments, mainly the instrumental ones, appear as demonstrations, which are

planned to be carried out by the students themselves.

When writing the lab manual, our main object was to enable students to work

independently in the laboratory, observe relevant phenomena, draw conclusions and

make notes, by which anyone can reproduce the experiment.

It is strongly recommended, that students carefully study Chapter 1,

"Laboratory Safety", before starting work in the laboratory to avoid accidents.

Chapter 2 ("Laboratory Utensils and Methods") and 3 ("Simple Laboratory

Separation Methods") comprise general laboratory methodology, which are to be

mastered so that it can provide a firm knowledge throughout the course. The

following "Chemical Analysis" section includes qualitative, quantitative, organic and

bio-organic analytical practices. In Chapter 5 ("Reaction Kinetics") and Chapter 6

("Chemical Equilibria") simple test-tube reactions help students comprehend the basic

concepts of the processes also important in life sciences. A separate chapter, number

7, is devoted to the most powerful separation method, chromatography, widely used in

biomedical practice and research. Chapter 8, "Instrumental Analysis", offers an insight

into the high-performance methods employed in modern diagnostic laboratories.

In Chapter 9 (Appendix) one can find an overview of Inorganic Chemistry,

particularly recommended in the first half of the course. Calculation Exercises are

attached to promote student's own preparation for the exams on General Chemistry.

Concentrations of reagent solutions used regularly in the practices, which are

prepared by technicians in the stock laboratory, are not indicated in the text. For a

better reproducibility, they are also available in the Appendix.

Our special appreciation is extended to Miss Andrea Gungl and Mrs Éva

Halász for the preparation of figures, and to Mrs Gabriella Németh for the computer-

edited structural formulas.

Despite multiple painstaking revisions, there must have some misprints and

other errors left in the manual. Notices on correction and improvement are thankfully

welcome.

21 May 1995

The Editor

3

1.1.1.1. INTRODUCTION INTRODUCTION INTRODUCTION INTRODUCTION

LABORATORY SAFETY

When working in a chemical laboratory we are handling several chemicals

with more or less adverse effects to human health, and we are performing experiments

that have a number of potential hazards associated with them. Thus, a chemical

laboratory can be a dangerous place to work in. With proper care and circumspection,

strictly following all precautionary measures, however, practically all accidents can be

prevented.

It is the prevention of accidents and damages posed by the speciality of the

chemical laboratory experiments that requires you to follow the instructor's advice as

well as keep the laboratory order during work in the laboratory. You should never

forget that your carelessness or negligence can threaten not only your own safety but

that of your classmates working around you.

This section has guidelines that are essential to perform your experiments in a

safe way without accident.

Preparation in advance

a) Read through the descriptions of the experiments carefully. If necessary, do

study the theoretical background of the experiments from your chemistry book(s).

After understanding, write down the outline of the experiments to be performed in

your laboratory notebook. If any items you don't understand remain, do ask your

instructor before starting work.

b) Prepare your notebook before the laboratory practice. Besides description of

the outline of the experiments, preliminary preparation should also include a list of the

chemicals and procedures which need special care and attention to avoid laboratory

accidents during your work.

Laboratory rules

a) The laboratory instructor is the first to enter and the last to leave the

laboratory. Before the instructor's arrival students must not enter the laboratory.

4

b) Always wear laboratory coat and shoes in the laboratory. Sandals and open-

toed shoes offer inadequate protection against spilled chemicals or broken glass.

c) Always maintain a disciplined attitude in the laboratory. Careless acts are

strictly prohibited. Most of the serious accidents are due to carelessness and

negligence.

d) Never undertake any unauthorized experiment or variations of those

described in the laboratory manual.

e) Maintain an orderly, clean laboratory desk and cabinet. Immediately clean

up all chemical spills from the bench and wipe them off the outer wall of the reagent

bottles with a dry cloth.

f) Smoking, drinking, or eating is not permitted during the laboratory practice.

Do not bring other belongings than your notebook, stationery, and laboratory manual

into the laboratory. Other properties should be placed into the locker at the corridor.

g) Be aware of your neighbour's activities. If necessary, warn them of improper

techniques or unsafe practices.

h) At the end of the lab, completely clean all glassware used in the

experiments, clear the laboratory bench of reagent bottles, glassware and equipment,

and clean it with a dry cloth. After putting back all your personal labware into your

cabinet, lock it carefully.

i) Always wash your hands with soap before leaving the laboratory.

Handling chemicals and glassware

a) At the beginning of the laboratory practices the instructor holds a short

introduction when all questions related to the experimental procedures can be

discussed.

b) Perform each experiment alone. During your work always keep your

laboratory notebook nearby in order to record the results of the experiments you

actually perform.

c) Handle all chemicals used in the experiments with great care. Never taste,

smell, or touch a chemical or solution unless specifically directed to do so.

d) Avoid direct contact with all chemicals. Hands contaminated with

potentially harmful chemicals may cause severe eye or skin irritations.

5

e) Reactions involving strong acids, strong bases, or chemicals with

unpleasant odour should be performed under the ventilating hood. If necessary, safety

glasses or goggles should be worn.

f) When checking the odour of a substance, be careful not to inhale very much

of the material. Never hold your nose directly over the container and inhale deeply.

g) When carrying out the experiments, first read the label on the bottle twice to

be sure of using the correct reagent. The wrong reagent can lead to accidents or

"inexplicable" results in your experiments.

h) Do not use a larger amount of reagents than the experiment calls for. Do not

return any reagent to a reagent bottle. There is always the chance that you accidentally

pour back some foreign substance which may react with the chemical in the bottle in

an explosive manner.

i) Do not insert your own pipette, glass rod, or spatula into the reagent bottles;

you may introduce impurities which could spoil the experiment for the person using

the stock reagent after you.

j) Mix reagents always slowly. Pour concentrated solutions slowly and

continuously stirring into water or into a less concentrated solutions. This is especially

important when diluting concentrated sulphuric acid.

k) Discard waste or excess chemicals as directed by your laboratory instructor.

The sink is not for the disposal of everything.

l) Using clean glassware is the basic requirement of any laboratory work.

Clean all glassware with a test-tube brush and a detergent, using tap water. Rinse first

with tap water and then with distilled water. If dry glassware is needed, dry the wet

one in drying oven, or rinse with acetone and air dry it.

ACCIDENT PROTECTION, FIRE PROTECTION AND FIRST AID

Accident and fire protection

a) Before starting the experiments make sure all the glassware are intact. Do

not use cracked or broken glassware. If a glassware breaks during the experiment, the

chemical spill and the glass splinters should be cleaned up immediately. Damaged

glassware should be replaced from the stock laboratory.

b) Fill not more than 4-5 cm3 of reagent into a test tube. As you are

performing the experiments, do not look into the mouth of the test tube and do not

6

point it at anyone. If you want to check the odour of a substance formed in a test tube

reaction, waft the vapours from the mouth of the test tube toward you with your hand.

c) Before heating glassware make sure that its outer wall is dry. Wet glassware

can easily break on heating. When heating liquids in a test tube, hold it with a piece of

tightly folded paper or a test-tube holder.

d) When heating liquids in an Erlenmeyer flask or in a beaker, support the

glassware on a wire gauze placed on an iron tripod, and put a piece of boiling stone

into the liquid to prevent bumping. Start heating with a law flame and intensify it

gradually.

e) When lighting the Bunsen burner, close the air-intake holes at the base of

the burner, open the gas cock of the outlet, and bring a lighted match to the mouth of

the burner tube until the escaping gas at the top ignites. (It is advantageous to strike

the mach first and then open the gas cock.) After it ignites, adjust the air control until

the flame is pale blue and the burner produces a slight buzzing sound.

f) If the Bunsen burner "burns in", which can be noticed from its green flame

and whistling (whizzing) sound, the gas valve of the outlet should be turned off

immediately. Allow the burner to cool, and light it again as described above.

g) When using an electric heater or other electric device, do not touch them

with wet hands and prevent liquids from spilling over them. If it accidentally happens

(e.g. a flask cracks on heating), unplug the device immediately and wipe off the liquid

with a dry cloth.

h) As a general rule, a flame should be used to heat only aqueous solutions.

Most organic solvents boil below the boiling point of water. A hot water bath can be

effectively used to heat these solvents.

i) When working with flammable organic solvents (e.g. hexane, diethyl ether,

petroleum ether, benzene) use of any open flame in the laboratory is prohibited. The

vapours of the flammable substances may waft for some distance down their source;

thus presenting fire danger practically in the whole laboratory.

j) In case of a smaller fire (e.g. a few millilitres of organic solvent burning in a

beaker or an Erlenmeyer flask), it can be extinguished by placing a watch glass over

the mouth of the flask. In case of a bigger fire and more serious danger, use the fire

extinguisher fixed on the wall of the laboratory. At the same time alarm the University

Fire Fighter Office by calling the N° 1333 from the corridor or from the stock lab.

k) Never blow the fire. This way you might turn the fire up and the flame can

shoot into your face. Do not use water to smother fires caused by water-immiscible

7

chemicals (e.g. benzene) and alkali metals. Pouring water on a plugged electric device

is also prohibited.

l) If your clothing catches fire, you can smother the flames by wrapping

yourself in a wet towel or a laboratory coat.

m) In case of fire in the laboratory the main gas valve and the electric switch

of the laboratory should be turned off immediately. (They are located in the corridor

on the outer wall of the laboratory.) Besides fighting the fire, start giving the injured

first aid immediately.

First aid

a) In case of an accident or injury, even if it is minor, notify your laboratory

instructor at once. The urgent first aid is an absolute requirement for the prevention of

more serious adverse health effects.

b) Minor burns caused by flames or contact with hot objects should be cooled

immediately by flooding the burned area with cold water, then treating it with an

ointment. Severe burns must be examined by a physician.

c) In case of a cut, remove the contamination and the glass splinters from the

wound. Disinfect its boundary with alcoholic iodine solution and bind it up with

sterile gauze. In case of severe cases the wound should be examined and treated by a

physician.

d) Whenever your skin gets into contact with chemicals, wash it quickly and

thoroughly with water. In case of chemical burns, the chemical should be neutralized.

For acid burns, the application of a dilute sodium hydrogen carbonate, for burns by

alkali, the application of a dilute solution of boric acid is used. After neutralization,

wash the affected area with water for 5-10 minutes and apply an oily ointment if

necessary.

e) Concentrated sulphuric acid dripped onto your skin must be wiped away

with a dry cloth. Then the affected area should be treated as described for acid burns

above.

f) Acids splashed onto your clothes could be neutralized with diluted solution

of ammonia or sodium hydrogen carbonate.

g) If any chemical gets into your mouth, spit it out immediately, and wash your

mouth well with water. In case of acidic or basic chemicals rinse your mouth first with

a diluted solution of sodium hydrogen carbonate or that of boric acid, respectively, to

neutralize the acid or the base.

8

h) If any chemical enters your eyes, immediately irrigate the eyes with large

quantities of water. In case of any kind of eye damage consult a physician

immediately.

i) In case of inhalation of toxic chemicals the injured should be taken out to

fresh air as soon as possible.

j) In case of an electric shock, the immediate cutoff of the electric current

supply of the laboratory (main switch) is the most important step to avoid irreversible

health damage. The injured should get medical attention as soon as possible. If

necessary, artificial respiration should be given until the physician arrives.

9

2. LABORATORY UTENSILS AND METHODS

UTENSILS

The most common laboratory utensils are shown in Fig. 2.1a-b

SIMPLE LABORATORY METHODS

Laboratory Measurements

Concerning basic laboratory measurements (weighing, measuring volume,

temperature, density, etc.) and data processing we refer to Biophysics and Biometrics.

Heating

In the laboratories, both direct and indirect methods of heating are used. For

direct heating a free gas flame or an electric hotplate or an electric heating mantle, for

indirect heating mostly liquid baths (water, mineral oil, silicone oil, glycerol) are used.

One of the most common sources of heat is the gas Bunsen burner (Fig.2.1b).

Solid materials can be heated (or ignited) on a direct flame in porcelain or

metal crucibles placed onto an iron tripod by means of a clay triangle. Heating is

started with a small flame and intensified gradually. Ignitions under precisely

controlled conditions can also be performed in an electric furnace.

Smaller amounts of aqueous solutions may be heated and boiled on direct

flame keeping a test tube with not more than 3-4 cm3 of the solution into the mantle

of the flame by means of a strip of a triple-folded paper sheet or a test-tube holder.

First the upper part of the solution is heated and then the rest. Caution: Never turn the

opening of a heated test tube toward your face or toward other people working

nearby.

Larger amounts of nonflammable, not too volatile liquids can be heated and

boiled in an Erlenmeyer flask or a beaker placed onto an iron tripod by means of an

10

asbestos wire gauze. In case of boiling, a piece of pumice (boiling stone) is placed into

the solution to avoid bumping.

Fig. 2.1a; Laboratory utensils

11

Fig.1b; Laboratory utensils

1. Test tube 2. Conical (Erlenmeyer) flask

3. Beaker 4. Wash bottle

5. Measuring cylinder 6. Glass funnel

7. Watchglass 8. Porcelain crucible

9. Petri dish 10. Wide-mouthed reagent bottle

11.Narrow-mouthed reagent bottle 12. Glass rod

13. Crucible tongue 14. Plastic spoon

15. Bunsen burner 16. Asbestos wire gauze

17. Clay triangle 18. Porcelain dish

12

Flammable liquids in a smaller amount can be heated on an electric hotplate

or in a water bath under the hood. Larger amounts of flammable liquids should be

heated and boiled in flasks equipped with a reflux condenser (Fig. 3.8) by a suitable

liquid bath or an electric heating mantle.

Forming Precipitates

Precipitates are produced mostly for analytical purposes. Then the visual observation

is very important. Therefore, the reagent is added gradually with shaking while the

changes are monitored. For a complete precipitation, the reagent should be added in

excess. Note that in case of amphotery or complex formation the excess of the reagent

can dissolve the previously formed precipitate.

Preparation of Gases

Laboratory gases (O2, N2, Cl2, H2, CO2, N2O, HCl, noble gases) on large scale are

put on the market in high-pressure cylinders. Smaller amounts of them can also be

prepared in the laboratory. The most widely used gas generator is the so-called Kipp's

apparatus which is suitable for producing gases from a solid, water-insoluble and a

liquid reactant (Fig. 2.2). It consists of two bulbs separated by a diaphragm and of a

long-stem funnel introducing the solution into the lower bulb. The apparatus functions

semi-automatically as follows:

Fig. 2.2; Kipp's apparatus

13

The solid reactant is placed into the middle bulb and, from the upper funnel,

the liquid reactant is introduced into the lowest bulb. The lowest bulb can not be filled

because of the resistance of the compressed air in the two lower bulbs. Turning the

gas-outlet stopcock on, the hydrostatic pressure pushes the liquid reactant through the

diaphragm into the middle part and the reaction starts. If the gas cock is turned off, the

increasing pressure of the evolving gas presses the liquid back into the lowest part and

the reaction is stopped. However, the apparatus is then ready to function again if the

cock is turned on. Usually a wash bottle is attached to the gas outlet for washing out

the drops of the reagents and for flow control. A rate of not more than 2 to 3 bubbles

per second gas flow is advised.

Gases frequently prepared in Kipp's apparatus:

Hydrogen from granulated zinc and hydrochloric acid according to the equation:

Zn + 2 H+ = H2 + Zn2+

Hydrogen sulphide from iron(II) sulphide and hydrochloric acid:

FeS + 2 H+ = H2S + Fe2+

Carbon dioxide from cracked marble and hydrochloric acid:

CaCO3 + 2 H+ = Ca2+ + H2O + CO2

14

3. SIMPLE LABORATORY SEPARATION METHODS

Decantation

The simplest way of the separation of a liquid from a solid phase is called decantation.

In this operation a suspension is allowed to stand until the precipitate settles down and

then the supernatant liquid is cautiously poured off. Then, the rest is suspended in

distilled water and the procedure is usually repeated twice to remove the dissolved and

adsorbed impurities. For the sedimentation of very fine suspensions, especially for

that of colloidal ones, centrifugation is used.

Filtration

Filtration means the separation of two phases with the aid of a filtering layer which

allows only one of the phases to pass through. Usually the solid phase is separated

from the liquid and rarely from the gaseous phase. In case of simple filtration,

hydrostatic pressure is enough to have the liquid pass through. To enhance the rate of

filtration, reduced pressure can be applied for vacuum filtration in the laboratory.

Rarely, filtration under pressure is also applied.

If simple filtration is performed at atmospheric pressure, smooth, conical and

short-stem funnels are used (Fig. 3.1).

Fig. 3.1; Glass funnels

15

Drying

Drying means the removal of an unnecessary liquid or its vapour from a solid, liquid

or gaseous material. In most cases, water is to be removed. The performance of drying

depends on the state of matter and the thermal stability of the substance to be dried.

Thus, gases and liquids are usually kept in direct contact with the drying agent. Solid

samples are usually placed into a closed space (desiccator) and the drying agent acts

through the common airspace (Fig. 3.6). A wide variety of hygroscopic materials are

used as dehydrating agents. Of hydrate-forming substances, anhydrous CaCl2,

Mg(ClO4)2, Na2SO4, K2CO3 cc. H2SO4, NaOH and KOH are widely used. Some

metals and a few oxides bind water by a chemical reaction, e.g. Na, K, CaO, P2O5.

Silica gel, which is coloured by a blue cobalt salt when prepared for drying, binds

water by adsorption. If silica is saturated with water, the blue colour of the cobalt salt

turns pink, indicating that silica should be reactivated by warming.

Fig. 3.6; Desiccator; infrared lamp

Solids, at room temperature, are dried in a desiccator (Fig 3.6). The

dehydrating agent is placed at the bottom of the apparatus, and the substance to be

dried onto the perforated plate. Reduced pressure may be applied to increase the

evaporation rate (vacuum desiccator). Thermally stable samples may be dried in an

electric oven or under an infrared lamp (Fig. 3.6).

16

Gases can be dried by passing them through an adsorption tube with a solid

packing, e.g. P2O5 or silica or through a washing bottle filled with a liquid drying

agent, e.g. cc. sulphuric acid.

Liquids are usually dried by adding an insoluble solid dehydrating agent and

are allowed to stand for a longer period of time. The drying agent can be removed by

filtration or decantation. For the dehydration of certain organic solvents, azeotropic

distillation (see later, at "Distillation") is useful.

Demonstration: Desiccator, infrared lamp.

17

4. CHEMICAL ANALYSIS

The aim of chemical analysis is the qualitative recognition and the determination of

the quantity of the constituents (e.g. atoms, ions, groups) of substances or mixtures.

Accordingly, analytical chemistry can be classified into two main parts: qualitative

and quantitative analysis. However, in the modern natural science and industry it is

also requested to answer questions about the form, distribution and interactions of the

constituents. Therefore, the task of analysis is to answer all the questions about the

chemical composition, structures, etc. being important in the characterization or

usage.

Chemical analysis may be classified according to the nature of the substances

investigated. Thus, inorganic and organic analysis are different but not strictly

independent fields of analysis. Inorganic analysis may be subdivided into metal,

silicate, etc. analysis; organic analysis comprises hydrocarbon, protein, food,

pharmaceutical, etc. analysis.

Another way of classification is based on the applied methods. The two main

branches of quantitative analysis are gravimetry and titrimetry. Gravimetry is based on

the measurement of the mass of the samples, while the volumetric methods are based

on the measurement of volumes proportional with the amount of a gas or a dissolved

reactant. Many of the physical or physicochemical properties are known to give

measurable values proportional with the mass or the concentration of different

samples. These instrumental analysis methods allow to perform the gravimetric or

volumetric determinations more precisely and sometimes can be automated. On the

other hand, instrumental methods themselves are useful to determine concentration or

composition occasionally together with structure. Separation methods represent

another enormously developing field of analysis.

It should be noted that the modern, fast and high performance methods unify

the separation, detection and quantitative determination techniques as well. The

principles are the following:

a) During separation the components are identified and usually quantitatively

measured. (E.g. immunoelectrophoresis, affinity chromatography or the combined gas

chromatography-mass spectrometry in which the chromatographically separated and

quantitatively measured components are identified by the attached mass

spectrometer.)

18

b) Without separation, selective analysis methods are used for one or more

components of a mixture not being sensitive to other components present (e.g.

enzymatic determination of blood glucose).

19


Qualitative analysis deals with the identification of elements and compounds

respectively, with the detection of individual components in their mixtures. Often

from the physical properties some conclusion may be drawn about the quality. Thus,

density, melting point, boiling point, colour, odour, etc. may carry partial information.

Since reliable conclusion, as regards the composition, can be drawn only from

chemical behaviour, the chemical reactions of a sample should be studied. Various

reagents should be added and the physical-chemical changes observed. The reactions

should be rapid, sensitive and selective.

Most of the reactions are carried out in aqueous solutions and the observable

changes are as follows:

a) The reagent reacts with one or more components of the sample forming an

insoluble precipitate. Further information can be drawn from the colour of the

precipitate and its behaviour against other reagents.

b) In other cases the reaction is accompanied with gas evolution. Then, the

physical-chemical properties of the gas may be informative.

c) The reagent brings about a colour-change in the solution.

d) Certain substances placed into gas flame change its colour (flame test).

Depending on the amounts of the samples and the expenses, test-tube, spot

and microchemical reactions can be carried out. Test-tube reactions are performed

with 1-2 cm3 of the sample in test tubes. The reagent is added dropwise with shaking,

occasionally with gentle heating. Spot reactions mean using 1-2 drops of the reactants

on a watch-glass, a porcelain dish or a filter paper. Microchemical reactions are

carried out under a microscope with very little amounts.

20

DETECTION OF CATIONS

During the systematic qualitative analysis the test for cations always precedes

that of anions since the previous knowledge of the present cations simplifies the

analysis of anions.

The most common cations are classified into five analytical groups taking

advantage of the different solubility of their derivatives formed with the reagents

added subsequently in the order: hydrochloric acid, hydrogen sulphide, ammonium

sulphide and ammonium carbonate. The formulas of physiologically important

(essential or poisonous) ions are underlined. They will be studied in the laboratory

course.

To Group 1 belong cations which form a precipitate with hydrogen sulphide

in nitric acid solution, and the sulphide is insoluble in ammonium sulphide. The

group is subdivided into two subgroups according to the solubility of their chlorides.

Members of Group 1a give a precipitate with hydrochloric acid: Ag+, Pb2+, Hg22+.

Chlorides of cations of Group 1b are soluble in water. To this group Hg2+, Cu2+,

Bi3+, Cd2+ belong.

Group 2 consists of cations forming insoluble sulphides with hydrogen

sulphide in nitric acid solution. However, these sulphides are soluble in ammonium

sulphide in form of thio salts: As3+, As5+, Sb3+, Sb5+, Sn2+, Sn4+.

Group-3 cations form insoluble sulphides only in neutral or slightly basic

medium with ammonium sulphide. These sulphides are soluble in dilute hydrochloric

acid: Co2+, Ni2+, Fe2+, Fe3+, Cr3+, Al3+, Zn2+, Mn2+.

Group-4 cations do not react with the above reagents. They form insoluble

carbonate precipitate with ammonium carbonate in a neutral medium: Ca2+, Sr2+,

Ba2+.

Group-5 cations have no group reactions. They have to be detected with

specific reactions: Mg2+, Na+, K+, NH4+, Li+, H+.

21

GROUP 1a

Reactions of mercury(I) ion (Hg22+)

For testing the reactions of mercury(I) ion, mercury(I) nitrate [Hg2(NO3)2]

stock solution is used.

Experiment 4.1.6 (Subgroup reaction)

a) Addition of HCl solution leads to the formation of white crystals of

mercury(I) chloride (calomel).

Hg22+ + 2 Cl- = Hg2Cl2

b) Adding NH4OH solution to the crystals, elementary mercury and mercury

amido chloride form (disproportionation) and the slurry turns grey.

Hg2Cl2 + NH4+ + 2 OH- = Hg

NH2

Cl + Hg + Cl- + 2 H2O

Experiment 4.1.7 (Group reaction)

On the action of H2S gas, mercury(I) ions disproportionate to form a black

precipitate of elementary mercury and mercury(II) sulphide. The precipitate is

insoluble in acids.

Hg22+ + S2- = Hg + HgS

Experiment 4.1.8

Adding NaOH solution, black mercury(I) oxide precipitates.

Hg22+ + 2 OH- = Hg2O + H2O

Experiment 4.1.9

Adding NH4OH solution, metallic mercury and basic mercury(II) amido

nitrate form as a black precipitate (disproportionation).

22

2 Hg22+ + NO3

- + NH4+ + 4 OH- = HgO Hg

NH2

NO3

+ 2 Hg + 3 H2O

GROUP 1b

Reactions of mercury(II) ion (Hg2+)

For testing the reactions of mercury(II) ion, mercury(II) chloride [HgCl2]



Introducing H2S gas into an acidic mercury(II) solution, a black precipitate of

mercury(II) sulphide deposits.

Hg2+ + S2-= HgS

Experiment 4.1.11

NaOH solution precipitates yellow mercury(II) oxide. Difference from

mercury(I) ion which turns black with NaOH.

Hg2+ + 2 OH- = HgO + H2O

Experiment 4.1.12

NH4OH solution precipitates white mercury(II) amido chloride.

Hg2+ + Cl- + NH4+ + 2 OH- = Hg

NH2

Cl + 2 H2O

Experiment 4.1.13

Potassium iodide (KI) solution gives a red precipitate of mercury(II) iodide.

The precipitate dissolves in excess reagent to form colourless, water-soluble

tetraiodomercurate(II) ion.

Hg2+ + 2 I- = HgI2

23

HgI2 + 2 I- = [HgI4]

2-

The alkaline solution of the complex salt is the so-called Nessler's reagent, which is

used for testing NH4+.

GROUP 2.

Reactions of arsenic(III) ion (As3+)

For testing the reactions of arsenic(III) ion, arsenic trioxide [As2O3] stock solution is

used.


From a solution acidified with HCl (!), H2S gas precipitates yellow arsenic(III)

sulphide ( if necessary, gentle warming is possible).

As2O3 + 6 H+ 2 As3+ + 3 H2O

2 As3+ + 3 S2- = As2S3

The precipitate dissolves in excess (NH4)2S solution forming a colourless

solution of ammonium thioarsenite. (Difference from Group 1.)

As2S3 + 3 S2- = 2 AsS3

3-

Experiment 4.1.20

Bettendorf's test:

Bettendorf's reagent, tin(II) chloride (SnCl2) in cc. HCl solution (caution, very

caustic!), reduces arsenic compounds to elementary arsenic as a black precipitate or a

mirror on the wall of the test tube.

2 As3+ + 3 Sn2+ = 2 As + 3 Sn4+

Procedure:

Excess volume of Bettendorf's reagent is added to the test solution and is

allowed to stand for a few minutes.

24

Experiment 4.1.21

Gutzeit's test:

Elementary zinc (in hydrochloric acid medium) reduces the soluble arsenic

compounds to arsine gas (AsH3). Allowing the gas to come into contact with a filter

paper wetted with a drop of concentrated silver nitrate solution, a yellow spot appears

which turns black when applying a drop of water onto it.

AsO33- + 9 H+ + 3 Zn = AsH3 + 3 Zn

2+ 3 H2O

AsH3 + 6 AgNO3 = Ag3As . 3 AgNO3 + 3 HNO3

Ag3As . 3 AgNO3 + 3 H2O = 6 Ag + H3AsO3 + 3 HNO3

Procedure:

To 1 cm3 As-containing solution a few cm3s of HCl, a few drops of CuSO4

and a piece of zinc metal are added. A wad of cotton wool is applied into the mouth of

the test tube as a filter and the tube is covered with a piece of filter paper wetted with

a drop of 50 % silver nitrate. A yellow spot develops within a few minutes which

turns black when a drop of water is added.

25

GROUP 3

Reactions of iron(II) ion (Fe2+)

For testing the reactions of iron(II) ion, iron(II) sulphate [FeSO4] stock

solution is used.


From a neutral solution, (NH4)2S solution precipitates black iron(II) sulphide

which is soluble in HCl.

Fe2+ + S2- = FeS

FeS + 2 H+ = Fe2+ + H2S

Experiment 4.1.35

NaOH solution precipitates greenish-white iron(II) hydroxide, which, allowed

to stand on air, is oxidized to brown iron(III) hydroxide.

Fe2+ + 2 OH- = Fe(OH)2

4 Fe(OH)2 + 2 H2O + O2 = 4 Fe(OH)3

Experiment 4.1.36

K3[Fe(CN)6], potassium-hexacyanoferrate(III) solution precipitates dark blue

iron(II) hexacyanoferrate(III).

3 Fe2+ + 2 [Fe(CN)6]3- = Fe3[Fe(CN)6]2

26

Reactions of iron(III) ion (Fe3+)

For testing the reactions of iron(III) ion, iron(III) chloride [FeCl3] stock

solution is used.


From a neutral solution, (NH4)2S solution precipitates black iron(II) sulphide

and elementary sulphur.

2 Fe3+ + S2- = 2 Fe2+ S

Fe2+ + S2- = FeS

HCl dissolves FeS, and the milky colloidal dispersion of sulphur can be seen.

Experiment 4.1.38

NaOH solution deposits reddish-brown, gelatinous iron(III) hydroxide.

Fe3+ + 3 OH- = Fe(OH)3

Experiment 4.1.39

K4[Fe(CN)6], potassium hexacyanoferrate(II) solution precipitates iron(III)

hexacyanoferrate(II) (Prussian blue). The reaction is specific for Fe3+ ion.

4 Fe3+ + 3 [Fe(CN)6]4- = Fe4[Fe(CN)6]3

Experiment 4.1.40

NH4SCN, ammonium thiocyanate (ammonium rodanide) solution in slightly

acidic medium produces blood-red coloration. Iron(III) thiocyanate formed may be

extracted (transferred into another phase) by shaking with diethyl ether.

Fe3+ + 3 SCN- = Fe(SCN)3

27

GROUP 4

Reactions of calcium ion (Ca2+)

For testing the reactions of calcium ion, calcium chloride [CaCl2] stock

solution is used.


(NH4)2CO3 or Na2CO3 solution in neutral or slightly alkaline medium

precipitates white crystals of calcium carbonate.

Ca2+ + CO32- = CaCO3

Experiment 4.1.42

(NH4)2(COO)2, ammonium oxalate solution in slightly basic, neutral or acetic

acid medium precipitates white Ca oxalate. This reaction is characteristic and very

sensitive to Ca2+ ion.

Ca2+ + (COO)22- = Ca(COO)2

Experiment 4.1.43

Fig. 4.1.1; Flame test

28

Flame test. The volatile Ca compounds colour the gas flame brick-red.

General procedure:

Into a porcelain crucible, 1 cm3 of the sample solution (or 0.1-0.5 g of a solid

sample) and a granule of elementary zinc are placed and the crucible is almost totally

filled with 20 % hydrochloric acid solution. The non-luminous flame of the Bunsen

burner held horizontally above the intensively bubbling solution changes its colour

(Fig. 4.1.1).

GROUP 5

Reactions of magnesium ion (Mg2+)

Magnesium ion belongs to Group 5, since, in the presence of other ammonium

salts, it does not form a precipitate with (NH4)2CO3.

For testing the reactions of magnesium ion, magnesium chloride [MgCl2]


Experiment 4.1.44

NaOH solution precipitates white, gelatinous magnesium hydroxide.

Mg2+ + 2 OH- = Mg(OH)2

In the presence of ammonium salts, due to the decrease of OH- concentration, the

precipitate does not form (see mass-action law; solubility product and Exps. 4.1.65

and 6.10).

Experiment 4.1.45

In the absence of other ammonium salts, (NH4)2CO3 or Na2CO3 solution

precipitates white crystals of basic magnesium carbonate.

4 Mg2+ + 4 CO32- + 4 H2O = Mg(OH)2

.3 MgCO3.3 H2O + CO2

29

Experiment 4.1.46

Magnesia mixture (see Exp. 4.1.65 and Exp. 6.10) forms a white precipitate

with sodium hydrogen phosphate (a characteristic reaction of both magnesium and

phosphate ions).

Mg2+ + NH4+ + PO43- = MgNH4PO4

Reactions of sodium ion (Na+)

For testing the reactions of sodium ion, sodium chloride [NaCl] stock solution

is used.

Experiment 4.1.47

Flame test: (For the performance see Exp. 4.1.43.) Sodium chloride vapour

colours the flame intense yellow (very sensitive for sodium).

Experiment 4.1.48

In a neutral or slightly alkaline medium potassium hexahydroxoantimonate(V)

solution precipitates white sodium hexahydroxoantimonate(V).

Na+ + [Sb(OH)6]- = Na[Sb(OH)6]

Reactions of potassium ion (K+)

For testing the reactions of potassium ion, potassium chloride [KCl] stock

solution is used.

Experiment 4.1.49

In a not too dilute solution sodium hydrogen tartrate precipitates potassium

hydrogen tartrate as white crystals.

30

COO-

CH

CH

COOH

OH

OH

COOK

CH

CH

COOH

OH

OH=K+

+

Experiment 4.1.50

Na3[Co(NO2)6], sodium hexanitritocobaltate(III) solution precipitates yellow

potassium hexanitritocobaltate(III). Note: The reagent solution is dark.

3 K+ + [Co(NO2)6]3- = K3[Co(NO2)6]

Experiment 4.1.51

Flame test: (For the performance see Exp. 4.1.43.) Potassium chloride vapour

colours the flame pale violet. Traces of sodium may obscure the colour, but viewed

through a cobalt glass the yellow colour of sodium can be filtered.

Reactions of ammonium ion (NH4+)

For testing the reactions of ammonium ion, ammonium chloride [NH4Cl]


Experiment 4.1.52

On addition of a strong base, e.g. NaOH, ammonia gas evolves. Stick a piece

of wet red litmus paper into the middle of the convex side of a watch glass. In another

watch glass mix 2-2 drops of ammonium chloride and sodium hydroxide solutions

and cover the watch glass with the first one. The blue colour of the litmus paper

indicates the evolution of ammonia gas. (Difference from potassium ion.)

NH4+ + OH- = NH3↑ + H2O

31

Experiment. 4.1.53

Na3[Co(NO2)6] solution, precipitates yellow ammonium hexanitrito-

cobaltate(III).

3 NH4+ + [Co(NO2)6]

3- = (NH4)3[Co(NO2)6]

Experiment 4.1.54

The most sensitive reaction of ammonium ion is the Nessler reaction. In

strongly basic medium potassium tetraiodomercurate(II), [HgI4]2- produces brown

precipitate or brownish-yellow coloration depending on the concentration of the test

solution. Basic mercury(II) amido iodide forms.

NH4+ + 2 [HgI4]

2- + 4 OH- = HgO HgNH2

I + 7 I- + 3 H2O

32

DETECTION OF ANIONS

Similarly to cations, anions are classified by testing with HNO3, BaCl2 or

Ba(NO3)2 and AgNO3 solutions. It should be noted that while the cations are

separable by means of the group reagents, anions cannot be so.

Group-1 anions react with strong acids to form gases or a precipitate: CO32-,

HCO3-, SO3

2-,S2O32-, S2-, SiO3

2-, ClO-.

Group-2 anions do not react with acids but react with BaCl2 or Ba(NO3)2 to

form a precipitate: SO42-, PO4

3-, BO33-, F-, IO3

-, BrO3-.

Group-3 anions give a precipitate with AgNO3: Cl-, I-, Br-, CN-, SCN-,

[Fe(CN)6]4-, [Fe(CN)6]

3-.

Group 4 consists of anions having no common reaction. These ions should be

detected individually: NO3-, NO2

-, ClO3-, OH-, CH3COO

-, (COO)22-.

33

GROUP 1

Reactions of carbonate ion (CO32-)

For testing the reactions of carbonate ion, sodium carbonate [Na2CO3] stock

solution is used.


Addition of acids causes evolution of carbon dioxide (see also Exp. 4.1.56).

CO32- + 2 H+ = H2O + CO2↑

Experiment 4.1.58

Barium-nitrate solution precipitates white barium carbonate. The precipitate

dissolves in dilute acids with the evolution of carbon dioxide gas. (Difference from

SO42- and PO4

3- ions.)

Ba2+ + CO32- = BaCO3

BaCO3 + 2 H+ = Ba2+ + CO2 + H2O

Experiment 4.1.59

The water-soluble carbonates hydrolyze making the solution considerably

basic. The solution dropped with phenolphthalein turns intense pink (difference from

soluble hydrogen carbonates).

CO32- + 2 H2O 2 OH- + H2CO3

34

Reactions of hydrogen carbonate ion (HCO3-)

For testing the reactions of hydrogen carbonate ion, sodium hydrogen

carbonate [NaHCO3] stock solution is used.


Acids produce evolution of carbon dioxide.

HCO3- + H+ = H2O + CO2↑

Experiment 4.1.61

Hydrogen carbonate ion is less basic than carbonate ion. Phenolphthalein

shows pale pink colour only.

HCO3- + H2O OH- + H2CO3

Reactions of hypochlorite ion (ClO-)

For testing the reactions of hypochlorite ion, sodium hypochlorite [NaOCl]



HCl generates Cl2 gas which can be detected by a wet iodo-starch paper held

to the opening of the test tube. Cl2 oxidizes iodide ions to I2, and the latter forms a

blue inclusion complex with starch.

Cl- + ClO- + 2 H+ = H2O + Cl2↑

2 I- + Cl2 = I2 + 2 Cl-

35

GROUP 2

Reaction of sulphate ion (SO42-)

Sulphuric acid stock solution is used for testing sulphate ions.


Barium nitrate solution precipitates powdery barium sulphate which is

insoluble in dilute acids (difference from carbonate and phosphate ions).

Ba2+ + SO42- = BaSO4

Reactions of phosphate ion (PO43-)

For tests of phosphate ion, disodium hydrogen phosphate [Na2HPO4] stock

solution is used.


Ba(NO3)2 solution in neutral or slightly alkaline medium precipitates white

barium hydrogen phosphate which is soluble in dilute acids (difference from sulphate

ion).

HPO42- + Ba2+ = BaHPO4

Experiment 4.1.65

Magnesia mixture in alkaline medium precipitates magnesium ammonium

phosphate as white crystals. The precipitate dissolves in acids (see also Exps. 4.1.46

and 6.10).

PO43- + Mg2+ + NH4

+ = MgNH4PO4

36

Experiment 4.1.66

In a solution acidified with HNO3, large excess of ammonium molybdate

precipitates yellow ammonium phosphomolybdate, (NH4)3[P(Mo3O10)4]. To 1-2

drops(!) of phosphate solution, 1 cm3 nitric acid is added, followed by a gradual

addition of ammonium molybdate until a yellow colour develops. In a few minutes,

yellow crystals separate.

4 H2[Mo3O10] + PO43- + 3NH4

+ (NH4)3[P(Mo3O10)4] + 4 H2O

GROUP 3

Reactions of chloride ion (Cl -)

For tests of chloride ion, sodium chloride stock solution is used.


AgNO3 solution precipitates white, curd-like, crystalline silver chloride.

Cl- + Ag+ = AgCl

The precipitate is divided into four test tubes, and the following tests are performed:

a) Adding nitric acid, the precipitate does not dissolve,

b) adding ammonia, the precipitate dissolves forming diamminesilver complex ion,

c) adding sodium thiosulphate, solution the precipitate dissolves in the form of

complex dithiosulphatoargentate ions,

d) allowed to stand, the white precipitate gradually turns grey due to the

photosensitivity of silver salts:

AgCl + 2 NH3 = [Ag(NH3)2]+ + Cl-

AgCl + 2 S2O32- = [Ag(S2O3)]

3- + Cl-

2 AgCl + hν = 2 Ag + Cl2

37

Reactions of iodide ion (I -)

For testing iodide ion, potassium iodide stock solution is used.


AgNO3 precipitates yellow powder of silver iodide.

I- + Ag+ = AgI

The precipitate is divided into four test tubes and the following tests are performed:

a) When adding nitric acid, the precipitate does not dissolve,

b) adding ammonia, the precipitate dissolves partly forming diamminesilver complex

ion,

c) adding sodium thiosulphate solution, the precipitate dissolves easily in the form of

complex dithiosulphatoargentate ions,

d) allowed to stand, the yellow precipitate turns gradually grey due to the

photosensitivity of silver salts:

AgI + 2 NH3 = [Ag(NH3)2]+ + I-

AgI + 2 S2O32- = [Ag(S2O3)2]

3- + I-

2 AgI + hν = 2 Ag + I2

Experiment 4.1.69

Chlorine water sets free elementary iodine from iodides.

2 I- + Cl2 = 2 Cl- + I2

The brown solution is divided into two test tubes and

a) starch solution is added to form a blue inclusion complex. The reaction is extremely

sensitive to iodine.

b) carbon tetrachloride is added and the mixture is shaken. In the water-insoluble,

oxygen-free organic solvent iodine shows violet colour.

38

Experiment 4.1.70

Bromine water acts similarly to chlorine.

2 I- + Br2 = 2 Br- + I2

GROUP 4

Reactions of nitrite ion (NO2-)

For the tests of nitrite ion, sodium nitrite [NaNO2] stock solution is used.

Experiment 4.1.71

A little NaNO2 solution is acidified with acetic acid, and FeSO4 solution is

added. The solution turns brown due to the formation of (unstable) nitroso iron

sulphate, [Fe(NO)SO4]. (Nitrate reacts similarly but only in presence of cc. H2SO4.)

Experiment 4.1.72

To 1-2 cm3 distilled water 1-2 drops of NaNO2 solution are added. Then,

1-1 cm3 of Griess-Ilosvay 1 and 2 reagent solutions in this order. In dilute nitrite

solution intense red colour develops. (At higher concentrations nitrite ion oxidizes the

red dye and the solution turns pale yellow.)

In acidic medium aromatic amines like sulphanilic acid (reagent 1) are

diazotized by nitrous acid and the diazonium salt formed reacts with other aromatics

like α-naphthylamine (reagent 2) in an azo-coupling reaction resulting a deeply

coloured dye.

NH2HOSO2 + HNO2 + H+

N NHOSO2

++ H2O 2

HOSO2 N N++ NH2 HOSO2 N N NH2 + H +

39

Reactions of nitrate ion (NO3-)

For the tests of nitrate ion, sodium nitrate [NaNO3] stock solution is used.

Experiment 4.1.73

To a little NaNO3 solution cc. H2SO4 is added carefully, the solution is

cooled under the tap and iron(II) sulphate solution is carefully layered on. A brown

ring forms known from the reactions of nitrite ion (see Exp. 4.1.71).

Experiment 4.1.74

On addition of Zn + HCl, nitrate ion can be reduced to nitrite and gives

positive Griess-Ilosvay reaction (see Exp. 4.1.72).

Procedure:

Repeat Exp. 4.1.72 with sodium nitrate solution, adding zinc powder to the

reaction mixture.

NO3- + Zn +H+ = NO2

- + Zn2+ +H2O

40

SIMPLE ANALYSIS OF CATIONS AND ANIONS

Experiment 4.1.76

Simple analysis of cations and anions

The simple analysis of cations and anions of medical importance – in cases

when only one cation and one anion are present – is performed as shown in Tables 1

and 2.

Procedure:

Prior to the systematic analysis it is strongly recommended that preliminary

tests be carried out with a small portion of the sample (colour, odour, crystal form,

etc.). Then, 0.5-1 g of the sample is dissolved in 30-40 cm3 water. One-third of this

stock solution is used to test for cations, another third for anion analysis, the rest is put

aside. It is advisable to check the pH of the stock solution before systematic analysis.

First the cations are identified, since their knowledge limits the number of the

possible counterions. E.g., if the sample is water-soluble and Hg22+ cation has been

found, the presence of Cl- and I- ions should be excluded.

After the identification of the ion according to Tables 1 and 2, all the

characteristic reactions of the suspected ion should be positive to accept the analysis

correct. The analysis process should be described in details in the notebook, including

visual observations and chemical equations.

41

Table 1; Simple analysis of cations

Sample

Reagent

Observation

Ion?

I

X (= stock solution) + HCl: white precipitate

Pb2+ or Hg2

2+

X + NaOH:

white precipitate black precipitate

Pb2+ Hg2

2+

II

solution I + H2S:

yellow precipitate *

As3+

black precipitate Hg2+ or Cu2+

X + NH4OH: white precipitate Hg2+

blue precipitate Cu2+

III

X + (NH4)2S:

white precipitate

Zn2+

green prec., soluble in HCl Cr3+

black precipitate Fe2+, Fe3+or Co2+

X + NaOH: greenish precipitate Fe2+

brown precipitate Fe3+

blue precipitate Co2+

IV X + Na2CO3:

X: flame test:

white precipitate brick-red negative

Ca2+ or Mg2+

Ca2+

Mg2+

V

If all reactions have been negative so far:

X: flame test: yellow Na+

pale violet K+

X + Nessler's reagent: brown precipitate NH4+

X + acid-base indicator: acidic H+

* In presence of oxidative ions (Fe3+or NO2

-), milky colloidal solution of sulphur may form.

42

Table 2; Simple analysis of anions

Sample

Reagent

Observation

Ion?

X

+ HCl

gas evolution

CO32-, HCO3

-

or ClO-

the gas CO2:

Cl2 (KI paper):

CO32- or HCO3

-

ClO-

X

+ Ba(NO3)2

a) prec. + HNO3

white precipitate:

dissolves with effervescence:

dissolves:

insoluble:

(CO32-) *,

SO42- or PO4

3-

(CO32-) *

PO43-

SO42-

X

+ HNO3 + AgNO3

precipitate:

Cl- or I-

a) prec. + NH4OH

b) prec.+Na2S2O3

soluble in a) and b):

soluble in b) only:

Cl-

I-

X

+ Griess-Ilosvay reagent

if negative for NO2-

addition of Zn + HCl

red colour:

red colour:

NO2-

NO3-

* If detection failed in Group 1 (solution too dilute).


43

SELF-TEST QUESTIONS

1. Why can the sulphides of the cations of the first two analytical groups be

precipitated with H2S in acidic medium?

2. Explain, why it is unsuccessful in the case of Group-3 cations and succeeds

when applying (NH4)2S.

3. What do you know about the temperature dependence of the solubility of

lead(II) chloride?

4. Explain the dissolution of lead(II) hydroxide in excess NaOH solution?

5. How can you distinguish between Hg2+ and Pb2+ ions?

6. Write equation for dissolution of copper(II) hydroxide precipitate in excess

ammonia solution.

7. Write equations for the reactions of potassium iodide with Hg2+, Pb2+ and

Cu2+ ions.

8. What is the common reagent of the fourth analytical group of cations?

10. How can Hg22+, Hg2+, Pb2+, Cu2+, Fe2+, Fe3+, Mg2+ and NH4

+ ions be

distinguished using a single reagent?

11. Propose a further reagent to distinguish Hg22+ and Hg2+ ions.

12. How could you prepare Nessler's reagent and which ion is detected with it?

13. Describe the essence of the Bettendorf-reaction.

14. Write down the reactions suitable to distinguish Fe2+ and Fe3+ ions.

15. During storage, iron(II) ions are easily oxidized to iron(III) ions by the

atmospheric oxygen. Propose a reaction to check the purity of an iron sulphate

sample.

16. What is the most sensitive reaction of Ca2+ ions?

17. Why is it impossible to precipitate Mg(OH)2 in the presence of ammonium

ions?

18. How can Ca2+ and Mg2+ ions be distinguished?

19. How can Na+ and K+ ions be distinguished?

20. How can K+ and NH4+ ions be distinguished?

21. Collect methods for detection of H+.


44

22. What are the common reagents for the Group 1, 2 and 3 anions?

23. How can you detect and explain the basicity of CO32- and HCO3

- solutions?

24. How could you detect OCl- ion?

25. How could you detect and distinguish SO42- and PO4

3- ions?

26. How to distinguish Cl- and I- ions?

27. How to distinguish NO2- and NO3

- ions? Describe the Griess-Ilosvay

reaction (see Organic Chemistry).

45

9. APPENDIX

9.1. INORGANIC CHEMISTRY

A short overview

INTRODUCTION

The aim of this compilation is to offer the students a supplement

a./ for recapitulation their recent knowledge having been acquired at the high school,

b./ for internalization additional practically interesting pieces of information about elements and

compouns, exceeding the limited subject of the laboratory practices and

c./ of short references to the biomedical importance of some inorganic substances.

It should be noted that this short overview comprises only supplementary material;

theoretical and analytical details can be found elsewhere!

Inorganic chemistry studies the structures and structure-property relationships of the

elements and their compounds except those containing carbon-hydrogen bonds. (The latters are

studied by Organic Chemistry). Considering that the chemical-physical properties of the elements

are mostly determined by their electronic structure (in close connection with their position in the

periodic table), it is plausible to discuss them according to the periodic law:

1./ Alkali metals and alkali earth metals (elements of the s-block except H and He);

2./ Main group metals (Sn, Pb, Bi);

3./ Transition metals (elements of the d-block);

4./ Metalloids (the lower left area of the p-block);

5./ Non-metals (elements of the p-block with high electronegativity values):

6./ Noble gases (the elements of the column VIII. of the periodic law).

46

CLASSIFICATION OF THE INORGANIC COMPOUNDS

The scientific classification of the inorganic compounds is based on their bonding systems

and crystal types: Ionic compounds usually form ionic crystals; covalent compounds are of

network or molecular crystals; the so-called intermetallic compounds exist in the form of a

metallic crystal lattice. This grouping allows a rather easy understanding of structure-property

relationships but exceeds our requirements. Considering that most of the inorganic compounds

are either hydrogen and/or oxygen compounds or salts, the following classification may be

useful:

Hydrogen compounds

Hydrogen forms so-called binary hydrides with almost all elements.

Ionic hydrides are formed with the elements of the lowest electronegativity values (s-

block elements) in which hydride anion (H-) exists (NaH, CaH2). These salt-like hydrides are

strong bases (proton acceptors, electron pair donors) and strong reducing agents (electron

donors).

The hydrides of the non-metallic elements (p-block elements) are covalent compounds

(volatile, of molecular crystal lattice). Their properties depend strongly on the partner atoms.

Thus, the hydrogen halides are polar, soluble in water, their solutions are acidic and their acid

strength increases with the atomic number of the halogen. Hydrides of the oxygen group elements

are slightly acidic (H2S), H2O is amphoteric. The elements of the nitrogen group form basic

hydrides (NH3, PH3), the carbon group hydrides are neutral. The hydrides of silicon, aluminium

and boron represent a transition between the covalent and salt-like ones.

Hydrides of the d-block elements (transition metals) are of metallic type in which

hydrogen atoms(!) occupy the space between the metal ions in the crystal lattice.

Oxygen compounds

The oxygen compounds are classified as oxides, hydroxides and oxoacids. Salts of

oxoacids will be discussed separately.

Oxides

a./ Basic oxides are the oxides of the elements of lowest electronegativity, reacting with

water to form the corresponding hydroxide bases, i.e. they are base anhydrides. Examples: The

47

oxides of the s-block elements (Na2O, K2O, CaO, BaO), a few of the p-block elements (Bi2O3)

and of the transition metals of lower oxidation state (FeO, MnO, CoO).

b./ Acidic oxides: To this group belong the oxides of the highly electronegative non-

metals and metalloids (SO2, SO3, P2O3, P2O5, As2O3, As2O5, CO2, SiO2, NO2, N2O3, N2O5,

etc.) and a few of the transition metals at higher oxidation state (CrO3, Mn2O7, V2O5, OsO4).

These oxides are acid anhydrides; They hydrolyze with water to form acids.

c./ Neutral oxides: This group comprises first of all H2O (amphoteric), NO, CO, N2O

(not being anhydrides) and the water-insoluble oxides ( Al2O3, Fe2O3).

Hydroxides and oxoacids

The basic oxides react with water to form bases, the acidic oxides to form oxoacids. In

aqueous solution, bases dissociate into hydroxide ion and a metal ion, the acids into hydrogen

(oxonium) ion and an oxo-anion. Acids and bases react to form water and a salt. The basic

strength of hydroxides increases with the atomic radius of the metal atom, the acid strength

increases with the oxidation number of the central atom. There is an intermediate group of

hydroxides (amphoteric hydroxides) exhibiting both weakly acidic and basic reactions depending

on the medium.

Acid-base character of X(OH)n compounds

Oxidation number of X

General formula

Basic Amphoteric Acidic

+1 XOH Li, Na, K, I Cl

+2 X(OH)2 Mg, Ca, Ba, Cr, Mn, Fe, Co, Ni, Cu

Be, Zn, Sn, Pb

+3 X(OH)3 Fe, Bi Al, Cr, As, Sb B, P

+4 X(OH)4

H2XO3

Pb, Sn C, Si, S, Se

+5 H3XO4

HXO3

As P, N, Cl, Br, I

48

+6 H2XO4 S, Se, Mo, Cr, Mn

+7 HXO4 Mn, Cl, I

A few examples:

Hydrolysis of a basic oxide: CaO + H2O = Ca(OH)2

Dissociation of the base: Ca(OH)2 = Ca2+ + 2 OH-

Hydrolysis of an acidic oxide: CO2 + H2O = H2CO3

Certain acidic oxides, depending on the stoichiometric ratio of added water, form different acids:

Adding maximum amount of water orthoacids, with less amounts pyro- and metaacids are

produced respectively, e.g.:

P2O5 + 3 H2O = 2 H3PO4 orthophosphoric acid

P2O5 + 2 H2O = H4P2O7 pyrophosphoric acid

P2O5 + H2O = 2 HPO3 metaphosphoric acid

There are acid anhydrides which disproportionate during hydrolysis. These are regarded as the

common anhydrides of two different acids, e.g.:

2 NO2 + H2O = HNO3 + HNO2

Dissociation of an acid: H2CO3 = 2 H+ + CO3

2-

Reaction of an acid and a base: Ca(OH)2 + H2CO3 = CaCO3 + 2 H2O

Basic oxides react with acids, acidic oxides with bases and basic oxides with acidic oxides,

respectively, also to form salts.:

Basic oxide with an acid: CaO + H2CO3 = CaCO3 + H2O

Acidic oxide with a base: CO2 + Ca(OH)2 = CaCO3 + H2O

Basic oxide with acidic oxide: CaO + CO2 = CaCO3

49

Salts

Salts are regarded as products of the acid-base (neutralization) reactions (i.e. they are not

oxygen containing compounds by all means. See, e.g. NH3 + HCl = NH4Cl). Salts can be

classified as follows:

a./ Normal salts: Salts resulted in a stoichiometric neutralization reaction:

2 NaOH + H2SO4 = Na2SO4 + 2 H2O

3 KOH + H3PO4 = K3PO4 + 3 H2O

b./ Acid salts: Salts formed by an incomplete neutralization of a polybasic acid. The

name: acid salt refers to the composition, not to their hydrolysis! (Their reactions in aqueous

solution are not acidic by all means!):

KOH + H3PO4 = KH2PO4 + H2O

2 KOH + H3PO4 = K2HPO4 + 2 H2O

NaOH + H2CO3 = NaHCO3 + H2O

KOH + H2SO3 = KHSO3 + H2O

c./ Base salts are products of a partial neutralization of a polyvalent (polyacidic) base.

The name: base salt refers to the composition, not to their hydrolysis! (Their aqueous solutions

are not basic by all means!):

Bi(OH)3 + HNO3 = Bi(OH)2NO3 + H2O

Sb(OH)3 + HCl = SbOCl + 2 H2O

d./ Mixed salts: Salts formed in a reaction of one base with two acids:

Ca(OH)2 + Cl2 = CaClOCl + H2O

50

e./ Double salts: Occasionally, from a solution of two different salts, uniform crystals are

obtained. The composition of these crystals is stoichiometric, according to the sum of the two

formulas and, when dissolved, dissociate into all ionic components:

K2SO4 + Al2(SO4)3 = 2 KAl(SO4)2 (alum)

When dissolved in water:

KAl(SO4)2 = K+ + Al3+ + 2 SO4

2-

or, e.g . (NH4)2SO4 +FeSO4 = (NH4)2Fe(SO4)2 (Mohr's salt)

When dissolved in water:

(NH4)2Fe(SO4)2 = 2 NH4+ + Fe2+ + 2 SO4

2-

f./ Complex salts are coordination compounds composed of an undissociable, so-called

complex ion and a dissociable counterion. The complex ion itself is composed of a central metal

ion surrounded by the so-called ligands, coordinatively bound to the central ion. The number of

the ligands is called coordination number.

AgCl + 2 NH3 = [Ag(NH3)2]Cl

When dissolved:

[Ag(NH3)2]Cl = [Ag(NH3)2]+ + Cl- (diamminesilver(I) chloride

Sodium dithiosulphatoargentate(I) behaves similarly in aqeous solution:

Na3[Ag(S2O3)2] = 3 Na+ + [Ag(S2O3)2]

3-

Thio compounds

In both inorganic and organic chemistry, compounds having divalent sulphur atom(s)

instead of oxygen atom(s) of a structurally analogous, known compound, are called thio

compounds. Thus, the thio analogue of water is hydrogen sulphide (H2S), those of oxides are the

sulphides (CO2 / CS2, CaO / CaS etc.), that of sulphuric acid is thiosulphuric acid (H2SO4 /

51

H2S2O3), that of arsenous acid is thioarsenous acid (H3AsO3 / H3AsS3), that of cyanic acid is

thiocyanic acid (rodanic acid) (HOCN / HSCN), etc..

Acyl groups

Acyl groups, as imaginary derivatives of oxoacids, can be obtained by the removal of the

hydroxyl groups (not hydroxide ions!) from the acid molecules. Thus, acyl groups have free

valence (-ies) useful to construct formulas of acid derivatives, e.g. acyl halides:

A few examples:

Oxoacid Acyl group Acyl chlorlide

nitrous acid: HNO2 nitrosyl group –NO nitrosyl chloride Cl–NO

nitric acid: HNO3 nitryl group –NO2 nitryl chloride Cl–NO2

sulphurous a.: H2SO3 sulphinyl =SO sulphinyl chloride Cl2SO

or thionyl group or thionyl chloride

sulphuric a.: H2SO4 sulphonyl =SO2 sulphonyl chloride Cl2SO2 or sulphuryl group or sulphuryl chloride

carbonic a.: H2CO3 carbonyl g. =CO carbonyl chloride COCl2 or phosgene

phosphoric a.: H3PO4 phosphoryl g. ≡PO phosphoryl chloride POCl3

or phosphorus oxychloride

52

THE NOMENCLATURE OF INORGANIC IONIC COMPOUNDS

Compounds can be described by their formulas and by their systematic, trivial or

pharmaceutical (Latin) names. In both formulas and names cation(s) precede(s) anion(s), e.g.

NaCl: sodium chloride; K2HPO4: dipotassium hydrogen phosphate.

If the cation may exist in different oxidation states, the oxidation number of it is given in

Roman numerals in brackets attached to the name of the cation, like Hg2Cl2: mercury(I) chloride

and HgCl2: mercury(II) chloride.

Anions of non-oxoacids have a name ending -ide, in Latin -atum, e.g.. Br-: bromide /

bromatum; CN–: cyanide / cyanatum; etc. Names of the anions of oxoacids usually end -ate / -

icum, e.g. CO32–: carbonate / carbonicum. If the central atom of the anion can exist in two

oxidation states the name of the most oxidized anion ends

-ate / -icum, the other ends -ite / -osum, e.g. SO42–: sulphate / sulfuricum, but SO3

2–: sulphite /

sulfurosum. In case of four different oxidation states the lowest oxidation state is referred by

hypo- / hypo- the highest one by per- / hyper- prefixes. (see the oxoacids of chlorine in Table 1.).

The formulas of the complex ions or neutral complexes are put in square brackets, those

of the ligands in brackets. Inside the formula of a complex ion, the atomic symbol of the central

atom or ion precedes the formula(s) of the ligands, e.g. [Pt(Cl)6]2–, [Ni(CO)4]. The names of the

ligands end -o except H2O: aqua, NH3: ammine and CO: carbonyl. The names of the central

atoms of the neutral complex molecules and cations are the usual ones, the names of the complex

anions end -ate. The oxidation state of the metal is put in brackets after the name, e.g. [Fe(CO)5]:

pentacarbonyliron(0) and [Fe(H2O)6]2+: hexaaquairon(II) ion, but [Fe(CN)6]

4–:

hexacyanoferrate(II) ion.

53

Table 1. Examples of naming inorganic compounds

Formula Name Latin name

( trivial name )

Anions ...........................ide .....................atum

F– fluoride fluoratum

Cl– chloride chloratum

Br– bromide bromatum

I– iodide iodatum

CN– cyanide cyanatum

OH– hydroxide hydroxydatum

O2– oxide oxydatum

S2– sulphide sulfuratum

H– hydride

hypo...........ite hypo..................osum

OCl– hypochlorite hypochlorosum

...................ite ....................osum

ClO2– chlorite chlorosum

NO2– nitrite nitrosum

SO32– sulphite sulfurosum

AsO33– arsenite arsenicosum

.......................ate ............................icum

ClO3– chlorate chloricum

BrO3– bromate bromicum

IO3– iodate iodicum

NO3– nitrate nitricum

CO32– carbonate carbonicum

SO42– sulphate sulfuricum

S2O32– thiosulphate thiosulfuricum

PO43– phosphate phosphoricum

AsO43– arsenate arsenicum

MnO42– manganate manganicum

OCN– cyanate cyanicum

SCN– thiocyanate (rodanide) thiocyanicum

per...............ate hyper................icum

MnO4– permanganate hypermanganicum

ClO4– perchlorate hyperchloricum

54

Acids Name Latin name

HCl hydrogen chloride acidum chloratum

HI hydrogen iodide

HCN hydrogen cyanide

H2S hydrogen sulphide

HOCl hypochlorous acid ac. hypochlorosum

HClO2 chlorous acid ac. chlorosum

HClO3 chloric acid ac. chloricum

HClO4 perchloric acid ac. hyperchloricum

HNO2 nitrous acid ac. nitrosum

HNO3 nitric acid ac. nitricum

H2SO3 sulphurous acid ac. sulfurosum

H2SO4 sulphuric acid ac. sulfuricum

H3PO3 orthophosphorous acid ac. phosphorosum

H3PO4 orthophosphoric acid ac. phosphoricum

H2CO3 carbonic acid ac. carbonicum

H2SiO3 silicic acid ac. silicicum

Bases

KOH potassium hydroxyde kalium hydroxydatum

Ca(OH)2 calcium hydroxyde calcium hydroxydatum

Fe(OH)2 iron(II) hydroxyde

Fe(OH)3 iron(III) hydroxyde

Salts

NaCl sodium chloride (table

salt)

natrium chloratum !!!

NaClO3 sodium chlorate natrium chloricum !!!

As2S3 arsenic(III) sulphide

As2S5 arsenic(V) sulphide

CaSO4 calcium sulphate

(gypsum)

calcium sulfuricum

MgSO4 magnesium sulphate magnesium sulfuricum

(bitter salt)

Na2SO4 sodium sulphate

(Glauber's salt)

natrium sulfuricum

CuSO4 copper(II) sulphate cuprum sulfuricum

FeSO4 iron(II) sulphate ferrosum sulfuricum

55

Fe2(SO4)3 iron(III) sulphate

Hg2Cl2 mercury(I)chloride

(calomel)

hydrargyrum chloratum

mite

HgCl2 mercury(II)chloride

(sublimate)

Formula Name Latin name

NaH2PO4 sodium dihydrogen

phosphate

Na2HPO4 disodium hydrogen

phosphate

Na3PO4 trisodium phosphate

Bi(OH)2NO3 bismuth dihydroxide

nitrate

bismuthum subnitricum

(basic bismuth nitrate)

Complexes

[Fe(CO)5] pentacarbonyliron(0)

[Ag(NH3)2]Cl diamminesilver(I)

chloride

Na3[Ag(S2O3)2] sodium

dithiosulphatoargentate(I)

K4[Fe(CN)6] potassium

hexacyanoferrate(II)

K3[Fe(CN)6] potassium

hexacyanoferrate(III)

K2[HgI4] potassium

tetraiodomercurate(II)

Fe3[Fe(CN)6]2 iron(II)

hexacyanoferrate(III)

Fe4[Fe(CN)6]3 iron(III)

hexacyanoferrate(II)

For the practice, one of the most important properties of compounds is their solubility.

The water-solubility of the most common inorganic salts is tabulated below.

56

Anion Solubility

Soluble in water

NO3– every nitrate is soluble

Cl– most of chlorides are soluble except: AgCl, Hg2Cl2, PbCl2

Br– most of bromides are soluble except: AgBr, Hg2Br2, HgBr2 and PbBr2

I– most of iodides are soluble except: AgI, Hg2I2, HgI2 and PbI2

SO42- most of sulphates are soluble except: CaSO4, SrSO4, BaSO4, PbSO4,

Hg2SO4, Ag2SO4

ClO3– every chlorate is soluble

C2H3O2– every acetate is soluble

Insoluble in water

S2– most of sulphides are insoluble except alkali-, alkali earth metal and ammonium-

sulphides

OH– most of hydroxides are insoluble except: alkali hydroxides and Ba(OH)2, Sr(OH)2,

Ca(OH)2

CO32– most of carbonates are insoluble except: alkali metal and ammonium carbonates

SO32– most of sulphites are insoluble except: alkali metal and ammonium sulphites

PO43– most of phosphates are insoluble except alkali metal and ammonium phosphates.

57

9.3 CONCENTRATIONS OF REAGENT SOLUTIONS

Concentrations of reagent solutions made by the technicians in the stock laboratory and not

indicated in the text can be seen below.

Acetic acid 12 %

Alcoholic β-naphthol 5 % Alum 5 % Ammonia 5 % Ammonia-ammonium-chloride buffer

0,8 % ammonium-chloride + 0,33 % ammonia

Ammonium acetate 8 % Ammonium carbonate 10 % Ammonium chloride 10 % Ammonium iron(III) sulphate 5 % Ammonium molybdate 5 %, in 4 % ammonia solution Ammonium oxalate 3 % Ammonium sulphide 5 % ammonia solution saturated with H2S gas Ammonium thiocyanate 0,90 % Arsenic trioxide 0.5 % (dissolves in hot water!) Barium chloride 5 % Barium nitrate 9 % Benzoic acid 0,12 % Bettendorf's reagent 10 % tin(II) chlorde in cc. HCl Bismuth chloride 4 % + cc. HCl until dissolves Bromine in carbon tetrachloride 5 % Bromine-water 1 % Calcium chloride 11 % Calcium hydroxide 5 % Chlorine-water distilled water saturated with chlorine gas Chromium(III) sulphate 6 % Cobalt(II) nitrate 10 % Copper(II) sulphate 4 % Dimethyl glyoxime 1 % in ethanol 2,4-Dinitrophenyl hydrazine HCl 2 %, in 40 % ethanol 3,5-Dinitrobenzoic acid 0,21 % Disodium hydrogen phosphate 7 % Eriochrome black T 1 %, in 70 % isopropanol Ethanolic silver nitrate 2 %

58

Fehling I 4 % copper(II) sulphate Fehling II 20 % K-Na-tartarate in 15 % NaOH solution Formaldehyde 38 % Fructose 5 % Glucose 10 % Glycine 1 % Glycogen 1 % Griess-Ilosvay I 1 % sulphanilic acid in 40 % acetic acid Griess-Ilosvay II 3 % α-naphthylamine in 30 % acetic acid Hydrochloric acid 5 % Hydroquinone 5 % Iodine 1 % +2 % KI Iron(II) sulphate 10 % Iron(III) chloride 8 % Lactose 5 % Lead(II) acetate 10 % Lead(II) nitrate 8 % Lecithin 0,2 % Magnesia mixture 10 %; see also Exp. 6.10! Magnesium chloride 10 % Magnesium sulphate 5 % Mercury(I) nitrate 3 %, + nitric acid until dissolves Mercury(II) chloride 1,5 %

Methanolic α-naphthol 1 % Methyl orange 0,10 % Methyl red 0,2 %, in 60 % ethanol Methylamine HCl 40 %, acidified with a little HCl

β-Naphthol 5 % Nessler-reagent 6 % Hg(II) chloride + 7,5 % KI in 20 % NaOH Nickel sulphate 5 % Ninhydrin 0,2 % in ethanol Nitric acid 20 % ortho-Toluidine 6 % + 0,15 % thiocarbamide in cc. acetic acid Phenol 5 % Phenolphtalein 1 % in 70 % ethanol Phenylhydrazine HCl 10 % Picric acid 5 % Potassium chloride 8 % Potassium chromate 3 % Potassium chromate (acidic) 3 %, in 20 % sulphuric acid

59

Potassium chromium(III) sulphate 5 % Potassium hexacyanoferrate(II) 2 % Potassium hexacyanoferrate(III) 2 % Potassium hexahydroxoantimonate(V)

2 %

Potassium iodide 2 % Potassium permanganate 0,3 % in 5 % sulphuric acid Protein 4-5 egg-whites in 1 dm3 water Rubeanic acid 0,1 %, in 65 % ethanol Saccharose 20 % Salicylic acid 1 % Silver nitrate 0.8 % Sodium acetate 11 % Sodium acetate, concentrated 32 % Sodium carbonate 5 % Sodium chloride 8 % Sodium dihydrogen phosphate 5 % Sodium hexanitritocobaltate(III) 40 %, in 3 % acetic acid Sodium hydrogen carbonate 9 % Sodium hydrogen tartarate 8 % Sodium hydroxide 5 % Sodium hypochlorite 2 % Sodium nitrate 5 % Sodium nitrite 5 % Sodium nitroprusside 1 %; unstable! Sodium thiosulphate 10 % Starch 1 %-os, + 0,1 % salicylic acid Sulphosalicylic acid 20 % Sulphuric acid 10 % Thymol blue 0,5 % in ethanol Thymolphthalein 0,1 %, in 50 % ethanol Trichloroacetic acid 0,16 % Trisodium phosphate 5 % Tryptophane 0,5 % Zinc sulphate 4 %

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