Department Of Biochemistry and Medical Chemistry Medical School, University of Pécs
Qualitative chemical analysis from
LABORATORY EXPERIMENTS IN MEDICAL CHEMISTRY
Edited by GYÖRGY OSZBACH
PÉCS, 1995
CONTENTS
1. INTRODUCTION
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Laboratory safety..................................................................................................... 3 Accident protection, fire protection, first aid .......................................................... 5 2. LABORATORY UTENSILS AND METHODS
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Utensils.................................................................................................................... 9 Simple laboratory methods...................................................................................... 9 Laboratory measurements .................................................................................. 9 Heating ............................................................................................................... 9 Forming precipitates ........................................................................................ 12 Preparation of gases......................................................................................... 12 3. SIMPLE LABORATORY SEPARATION METHODS
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Decantation............................................................................................................ 14 Filtration ................................................................................................................ 14 Drying.................................................................................................................... 15 4. CHEMICAL ANALYSIS
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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
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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 computeredited 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
1. 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.
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b) Always wear laboratory coat and shoes in the laboratory. Sandals and opentoed 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.
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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
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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
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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. 7
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.
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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
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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 10
1. Test tube
Fig.1b; Laboratory utensils 2. Conical (Erlenmeyer) flask
3. Beaker
4. Wash bottle
5. Measuring cylinder
6. Glass funnel
7. Watchglass
8. Porcelain crucible
9. Petri dish 11.Narrow-mouthed reagent bottle
10. Wide-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
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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
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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
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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
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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). 15
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.
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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 spectrometer.)
components are identified by the attached mass
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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).
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4.1 QUALITATIVE CHEMICAL ANALYSIS
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.
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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+.
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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
Cl NH2
+ 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).
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2 Hg22+ + NO3- + NH4+ + 4 OH- = HgO Hg
NH2 NO 3
+ 2 Hg + 3 H2O
GROUP 1b
Reactions of mercury(II) ion (Hg2+)
For testing the reactions of mercury(II) ion, mercury(II) chloride [HgCl2] stock solution is used.
Experiment 4.1.10 (Group reaction) 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 H2 O
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
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HgI2 + 2 I- = [HgI4]2The 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.
Experiment 4.1.19 (Group reaction) 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 AsS33-
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.
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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 Zn2+ 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.
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GROUP 3
Reactions of iron(II) ion (Fe2+)
For testing the reactions of iron(II) ion, iron(II) sulphate [FeSO4] stock solution is used.
Experiment 4.1.34 (Group reaction) 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
25
Reactions of iron(III) ion (Fe3+)
For testing the reactions of iron(III) ion, iron(III) chloride [FeCl3] stock solution is used.
Experiment 4.1.37 (Group reaction) 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
26
GROUP 4
Reactions of calcium ion (Ca2+)
For testing the reactions of calcium ion, calcium chloride [CaCl2] stock solution is used.
Experiment 4.1.41 (Group reaction) (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
27
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] stock solution is used.
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
28
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.
29
COO K+ +
COOK
CH OH CH OH COOH
=
CH OH CH OH COOH
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] stock solution is used.
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
30
Experiment. 4.1.53 Na3[Co(NO2)6] solution, precipitates yellow ammonium cobaltate(III). 3 NH4+ + [Co(NO2)6]3- = (NH4)3[Co(NO2)6]
hexanitrito-
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 Hg
31
NH2 I
+ 7 I- + 3 H2O
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-, SO32-,S2O32-, S2-, SiO32-, ClO-. Group-2 anions do not react with acids but react with BaCl2 or Ba(NO3)2 to form a precipitate: SO42-, PO43-, 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-.
32
GROUP 1
Reactions of carbonate ion (CO32-)
For testing the reactions of carbonate ion, sodium carbonate [Na2CO3] stock solution is used.
Experiment 4.1.57 (Group reaction) 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 PO43- 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
33
Reactions of hydrogen carbonate ion (HCO3-)
For testing the reactions of hydrogen carbonate ion, sodium hydrogen carbonate [NaHCO3] stock solution is used.
Experiment 4.1.60 (Group reaction) 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] stock solution is used.
Experiment 4.1.62 (Group reaction) 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-
34
GROUP 2
Reaction of sulphate ion (SO42-)
Sulphuric acid stock solution is used for testing sulphate ions.
Experiment 4.1.63 (Group reaction) 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.
Experiment 4.1.64 (Group reaction) 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
35
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.
Experiment 4.1.67 (Group reaction) 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]+ + ClAgCl + 2 S2O32- = [Ag(S2O3)]3- + Cl2 AgCl + hν = 2 Ag + Cl2
36
Reactions of iodide ion (I -)
For testing iodide ion, potassium iodide stock solution is used.
Experiment 4.1.68 (Group reaction) 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]+ + IAgI + 2 S2O32- = [Ag(S2O3)2]3- + I2 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.
37
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.
HOS O2
+ NH2 + HNO 2 + H
N N
HOS O2
+ HOSO 2
N N
+
HOSO 2
NH2
38
N N
+ + 2 H2O
NH2 + H +
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
39
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.
40
Table 1; Simple analysis of cations Sample
I
II
Reagent
Observation
Ion?
X (= stock solution) + HCl:
white precipitate
Pb2+ or Hg22+
X + NaOH:
white precipitate black precipitate
Pb2+ Hg22+
solution I + H2S:
yellow precipitate *
As3+
X + NH4OH:
III
X + (NH4)2S:
black precipitate
Hg2+ or Cu2+
white precipitate
Hg2+
blue precipitate
Cu2+
white precipitate
Zn2+
green prec., soluble in HCl
Cr3+
black precipitate greenish precipitate
Fe2+
brown precipitate
Fe3+
blue precipitate
Co2+
X + Na2CO3:
white precipitate
Ca2+ or Mg2+
X: flame test:
brick-red
Ca2+
negative
Mg2+
yellow
Na+
pale violet
K+
X + NaOH:
IV
V
Fe2+, Fe3+or Co2+
If all reactions have been negative so far: X: flame test:
X + Nessler's reagent:
brown precipitate
X + acid-base indicator:
acidic
NH4+ H+
* In presence of oxidative ions (Fe3+or NO2-), milky colloidal solution of sulphur may form.
41
Table 2; Simple analysis of anions Sample
X
Reagent
+ HCl
Observation
Ion?
CO32-, HCO3-
gas evolution
or ClO-
the gas
CO2:
CO32- or HCO3ClO-
Cl2 (KI paper):
X
+ Ba(NO3)2
white precipitate:
(CO32-) *, SO 2- or PO 34
a) prec. + HNO3
X
X
+ HNO3 + AgNO3
dissolves:
(CO32-) * PO 3-
insoluble:
SO42-
precipitate:
Cl- or I-
dissolves with effervescence:
a) prec. + NH4OH
soluble in a) and b):
b) prec.+Na2S2O3
soluble in b) only:
+ Griess-Ilosvay reagent
4
ClI-
red colour:
NO2-
red colour:
NO3-
if negative for NO2addition of Zn + HCl
4
* If detection failed in Group 1 (solution too dilute).
42
4.1 QUALITATIVE CHEMICAL ANALYSIS
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+.
43
4.1 QUALITATIVE CHEMICAL ANALYSIS
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 PO43- 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).
44
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 (sblock 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 nonmetals 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
Pb, Sn
C, Si, S, Se
As
P, N, Cl, Br, I
H2XO3 +5
H3XO4 HXO3
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+
Hydrolysis of an acidic oxide:
=
CO2
+
+
H2O
2 OH=
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
=
metaphosphoric acid
2 HPO3
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
Dissociation of an acid:
H2CO3
Reaction of an acid and a base:
Ca(OH)2
+ =
HNO2 2 H+
CO32-
+
+ 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
Acidic oxide with a base:
CO2
Basic oxide with acidic oxide:
+
= CaCO3
Ca(OH)2
CaO + CO2
+ H2O
= CaCO3 =
CaCO3
+ H2O
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
Sb(OH)3
+
HCl
=
SbOCl
+
H2O
+ 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:
or, e.g .
KAl(SO4)2
=
(NH4)2SO4
+FeSO4
K+
=
+
Al3+
+
2 SO42-
(NH4)2Fe(SO4)2 (Mohr's salt)
When dissolved in water: (NH4)2Fe(SO4)2
2 NH4+
=
+
Fe2+
+ 2 SO42-
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:
–NO2
nitryl chloride
Cl–NO2
=SO
sulphinyl chloride
Cl2SO
HNO3 nitryl group
sulphurous a.: H2SO3 sulphinyl
or thionyl group
or thionyl chloride
sulphuric a.: H2SO4 sulphonyl =SO2 or sulphuryl group
sulphonyl chloride
carbonic a.: H2CO3 carbonyl g.
carbonyl chloride
Cl2SO2
or sulphuryl chloride
=CO
COCl2
or phosgene phosphoric a.: H3PO4
phosphoryl g.
≡PO
phosphoryl chloride or phosphorus oxychloride
POCl3
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 2– 2– -ate / -icum, the other ends -ite / -osum, e.g. SO4 : sulphate / sulfuricum, but SO3 : 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]: hexaaquairon(II) ion, but [Fe(CN)6]4–: pentacarbonyliron(0) and [Fe(H2O)6]2+: hexacyanoferrate(II) ion.
53 Table 1. Examples of naming inorganic compounds
–
F Cl– Br–
Formula
Name
Latin name
Anions
( trivial name ) ...........................ide
.....................atum
fluoride chloride bromide
fluoratum chloratum bromatum
iodide
iodatum
CN OH–
cyanide hydroxide
cyanatum hydroxydatum
O 2– S 2–
oxide sulphide
oxydatum sulfuratum
H–
hydride
I– –
hypo...........ite hypochlorite ...................ite
hypo..................osum hypochlorosum ....................osum
ClO2– NO2–
chlorite nitrite
chlorosum nitrosum
SO32–
sulphite
sulfurosum
arsenite .......................ate
arsenicosum ............................icum
ClO3– BrO3–
chlorate bromate
chloricum bromicum
IO3– NO3–
iodate nitrate
iodicum nitricum
CO32– SO42–
carbonate sulphate
carbonicum sulfuricum
S2O32–
thiosulphate
thiosulfuricum
PO4 AsO43–
phosphate arsenate
phosphoricum arsenicum
MnO42– OCN– SCN–
manganate cyanate thiocyanate (rodanide)
manganicum cyanicum thiocyanicum
OCl
–
AsO3
3–
3–
per...............ate MnO4 ClO4–
–
permanganate perchlorate
hyper................icum hypermanganicum hyperchloricum
54 Acids HCl
Name hydrogen chloride
Latin name acidum chloratum
HI HCN
hydrogen iodide hydrogen cyanide
H 2S HOCl HClO2
hydrogen sulphide hypochlorous acid chlorous acid
ac. hypochlorosum ac. chlorosum
HClO3
chloric acid
ac. chloricum
HClO4 HNO2
perchloric acid nitrous acid
ac. hyperchloricum ac. nitrosum
HNO3 H2SO3
nitric acid sulphurous acid
ac. nitricum ac. sulfurosum
H2SO4
sulphuric acid
ac. sulfuricum
H3PO3 H3PO4 H2CO3
orthophosphorous acid orthophosphoric acid carbonic acid
ac. phosphorosum ac. phosphoricum ac. carbonicum
H2SiO3
silicic acid
ac. silicicum
KOH
potassium hydroxyde
kalium hydroxydatum
Ca(OH)2 Fe(OH)2
calcium hydroxyde iron(II) hydroxyde
calcium hydroxydatum
Fe(OH)3
iron(III) hydroxyde
Bases
Salts NaCl
sodium chloride (table
natrium chloratum !!!
salt) NaClO3 As2S3
sodium chlorate arsenic(III) sulphide
As2S5
arsenic(V) sulphide
CaSO4
calcium sulphate
natrium chloricum !!!
calcium sulfuricum
(gypsum) MgSO4 Na2SO4
magnesium sulphate (bitter salt) sodium sulphate
magnesium sulfuricum natrium sulfuricum
(Glauber's salt) CuSO4 FeSO4
copper(II) sulphate iron(II) sulphate
cuprum sulfuricum ferrosum sulfuricum
55 Fe2(SO4)3 Hg2Cl2 HgCl2
iron(III) sulphate mercury(I)chloride
hydrargyrum chloratum
(calomel)
mite
mercury(II)chloride (sublimate)
Formula NaH2PO4
Name sodium dihydrogen
Latin name
phosphate Na2HPO4
disodium hydrogen phosphate
Na3PO4 Bi(OH)2NO3
trisodium phosphate bismuth dihydroxide
bismuthum subnitricum
nitrate (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– Cl–
every nitrate is soluble most of chlorides are soluble except: AgCl, Hg2Cl2, PbCl2
Br–
most of bromides are soluble except: AgBr, Hg2Br2, HgBr2 and PbBr2 most of iodides are soluble except: AgI, Hg2I2, HgI2 and PbI2
–
I SO42-
most of sulphates are soluble except: CaSO4, SrSO4, BaSO4, PbSO4, Hg2SO4, Ag2SO4
ClO3– C2H3O2
every chlorate is soluble –
every acetate is soluble Insoluble in water
S2–
most of sulphides are insoluble except alkali-, alkali earth metal and ammoniumsulphides –
OH
most of hydroxides are insoluble except: alkali hydroxides and Ba(OH)2, Sr(OH)2, Ca(OH)2
CO3
2–
2–
SO3 PO43–
most of carbonates are insoluble except: alkali metal and ammonium carbonates most of sulphites are insoluble except: alkali metal and ammonium sulphites 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 Alcoholic β-naphthol Alum Ammonia Ammonia-ammonium-chloride buffer Ammonium acetate Ammonium carbonate Ammonium chloride Ammonium iron(III) sulphate Ammonium molybdate Ammonium oxalate Ammonium sulphide Ammonium thiocyanate Arsenic trioxide Barium chloride Barium nitrate Benzoic acid Bettendorf's reagent Bismuth chloride Bromine in carbon tetrachloride Bromine-water Calcium chloride Calcium hydroxide Chlorine-water Chromium(III) sulphate Cobalt(II) nitrate Copper(II) sulphate Dimethyl glyoxime 2,4-Dinitrophenyl hydrazine HCl 3,5-Dinitrobenzoic acid Disodium hydrogen phosphate Eriochrome black T Ethanolic silver nitrate
12 % 5% 5% 5% 0,8 % ammonium-chloride + 0,33 % ammonia 8% 10 % 10 % 5% 5 %, in 4 % ammonia solution 3% 5 % ammonia solution saturated with H2S gas 0,90 % 0.5 % (dissolves in hot water!) 5% 9% 0,12 % 10 % tin(II) chlorde in cc. HCl 4 % + cc. HCl until dissolves 5% 1% 11 % 5% distilled water saturated with chlorine gas 6% 10 % 4% 1 % in ethanol 2 %, in 40 % ethanol 0,21 % 7% 1 %, in 70 % isopropanol 2%
58 Fehling I Fehling II Formaldehyde Fructose Glucose Glycine Glycogen Griess-Ilosvay I Griess-Ilosvay II Hydrochloric acid Hydroquinone Iodine Iron(II) sulphate Iron(III) chloride Lactose Lead(II) acetate Lead(II) nitrate Lecithin Magnesia mixture Magnesium chloride Magnesium sulphate Mercury(I) nitrate Mercury(II) chloride Methanolic α-naphthol Methyl orange Methyl red Methylamine HCl β-Naphthol Nessler-reagent Nickel sulphate Ninhydrin Nitric acid ortho-Toluidine Phenol Phenolphtalein Phenylhydrazine HCl Picric acid Potassium chloride Potassium chromate Potassium chromate (acidic)
4 % copper(II) sulphate 20 % K-Na-tartarate in 15 % NaOH solution 38 % 5% 10 % 1% 1% 1 % sulphanilic acid in 40 % acetic acid 3 % α-naphthylamine in 30 % acetic acid 5% 5% 1 % +2 % KI 10 % 8% 5% 10 % 8% 0,2 % 10 %; see also Exp. 6.10! 10 % 5% 3 %, + nitric acid until dissolves 1,5 % 1% 0,10 % 0,2 %, in 60 % ethanol 40 %, acidified with a little HCl 5% 6 % Hg(II) chloride + 7,5 % KI in 20 % NaOH 5% 0,2 % in ethanol 20 % 6 % + 0,15 % thiocarbamide in cc. acetic acid 5% 1 % in 70 % ethanol 10 % 5% 8% 3% 3 %, in 20 % sulphuric acid
59 Potassium chromium(III) sulphate Potassium hexacyanoferrate(II) Potassium hexacyanoferrate(III) Potassium hexahydroxoantimonate(V) Potassium iodide Potassium permanganate Protein Rubeanic acid Saccharose Salicylic acid Silver nitrate Sodium acetate Sodium acetate, concentrated Sodium carbonate Sodium chloride Sodium dihydrogen phosphate Sodium hexanitritocobaltate(III) Sodium hydrogen carbonate Sodium hydrogen tartarate Sodium hydroxide Sodium hypochlorite Sodium nitrate Sodium nitrite Sodium nitroprusside Sodium thiosulphate Starch Sulphosalicylic acid Sulphuric acid Thymol blue Thymolphthalein Trichloroacetic acid Trisodium phosphate Tryptophane Zinc sulphate
5% 2% 2% 2% 2% 0,3 % in 5 % sulphuric acid 4-5 egg-whites in 1 dm3 water 0,1 %, in 65 % ethanol 20 % 1% 0.8 % 11 % 32 % 5% 8% 5% 40 %, in 3 % acetic acid 9% 8% 5% 2% 5% 5% 1 %; unstable! 10 % 1 %-os, + 0,1 % salicylic acid 20 % 10 % 0,5 % in ethanol 0,1 %, in 50 % ethanol 0,16 % 5% 0,5 % 4%
