Qualitative reactions to nitrogen. Download the book "Analytical chemistry of nitrogen" (2.24Mb)

ABOUT DETECTION

Methods of qualitative analysis, which do not require a lot of time, reagents and analyzed material to perform them, allow the analyst to quickly and easily assess the limits of the content of the element being determined (for the purpose of subsequent selection quantitative method its definition), and also in some cases obtain information about the forms of its presence in the sample under study.

Learn more about qualitative detection methods inorganic compounds nitrogen can be found in a number of manuals. Methods for detecting nitrogen in organic materials (organic qualitative analysis) are described in detail in the book. It also describes methods for converting total nitrogen into easily measurable forms. The work is devoted to the systematic microidentification of organic compounds, including methods for the rapid discovery of nitrogen using a Weiss ring bath (along with other important heteroatoms). The opening minimum of nitrogen is 0.01-1 mcg. Qualitative elemental analysis organic substances without their preliminary mineralization is described in the work. The work is devoted to the ultramicrocascillary method for the discovery of nitrogen in organic substances.

For the qualitative detection of nitrogen-containing ions, their specific chemical and physical properties are used: color reactions in test tubes, drop reactions, including on paper, microcrystalscopic reactions, sorption on AlO3, electrophoresis on paper, IR spectroscopy, fluorescence, catalytic methods, etc. d.

Below is short description The most common methods for the discovery of ammonium ions are nitrate, nitrite, thiocyanate and cyanide ions.

Caustic alkalis (NaOH, KOH) release gaseous ammonia from solutions of ammonium salts when heated, which is detected by smell using litmus or phenolphthalein paper.

Nessler's reagent, which is a mixture of complex salt K2 with KOH, forms a characteristic red-brown precipitate with solutions of ammonium salts (or, in very small quantities, a yellow color). The sensitivity of the reaction is 0.0003 mg in a drop of 0.002 ml. Ions of the elements Ag, Hg(II), Pb, and S2- ion interfere.

Nitrite ions

The acid decomposes all nitrites to form gaseous NO2, colored brown.

Potassium iodide in the presence of H2S04 is oxidized by nitrites to free Ja (other oxidizing agents also act: MnO4, CrOG, As04~).

An acetic acid solution of benzidine in the presence of NOj ions forms a yellow-colored compound.

Sulfanilic acid and a-naphthylamine (Griess-Idosvay reagent) in an acetic acid medium form a colored azo dye with nitrite ions.

A microcrystalloscopic reaction has been proposed for the detection of nitrite ions: a grain of the test substance is added to a drop of a solution containing potassium, lead and copper(II) acetate and acidified CH3COOH. Black K2Pb crystals stand out. This method allows you to open up to 0.75 mg of NOa - Limit dilution 1: 1500. The presence of NO3 ions does not interfere with the reaction.

K3 formation reaction. When the test solution is mixed with solutions of Co(N03)2, dilute acetic acid and KC1 in the presence of NO2, a yellow crystalline precipitate appears.

Potassium permanganate becomes discolored in an acidic environment when heated in the presence of nitrate ions as a result of the reduction of manganese to MPa+.

o-aminoanilide of benzenesulfonic acid (sulfuric acid solution) in an acidic environment precipitates NO2 ions.

Nitrate ions

Oxidation-reduction reactions are predominantly used to open nitrate ions.

Reaction with copper and sulfuric acid when heated leads to the release of brown gas NO2.

The reaction with FeS04 in the presence of concentrated HjS04 leads to the formation of a brown ring in a test tube as a result of the formation of the complex compound lFe(N03)]S04. Ions J -, Br~, oxidizing anions, SCN- interfere.

2 V. F. Volynets, M. P. Volynets

The reduction reaction to ammonia in the presence of a concentrated alkali solution with zinc dust, aluminum powder or Devard's alloy. NH3 is detected with litmus (blue) or phenolphthalein (red) paper. NHj", NOs, SCN", 2_ interfere. MnOj, SIO\~, N02 interfere.

Reaction of NO3 reduction to NO^ upon the action of zinc metal nitrates in the presence of CH3CO0H. Further, NO^ is detected by its characteristic reactions (see above).

Reaction with diphenylamine (G,Hs)aNH. Place 4-5 drops of a solution of diphenylamine in concentrated H2S04 on a watch glass. Add a little of the solution to be analyzed at the tip of a clean glass rod and mix. In the presence of NO3, an intense blue color appears due to the oxidation of diphenylamine by the resulting nitric acid. They interfere with N0^. Cr04~, Mn04, 3_, Fe3+ and other oxidizing agents,

Nitrogen- element of the 2nd period of the V A-group of the Periodic Table, serial number 7. Electronic formula of the atom [ 2 He]2s 2 2p 3, characteristic oxidation states 0, -3, +3 and +5, less often +2 and +4 and other state N v is considered relatively stable.

Scale of oxidation states for nitrogen:
+5 - N 2 O 5, NO 3, NaNO 3, AgNO 3

3 – N 2 O 3, NO 2, HNO 2, NaNO 2, NF 3

3 - NH 3, NH 4, NH 3 * H 2 O, NH 2 Cl, Li 3 N, Cl 3 N.

Nitrogen has a high electronegativity (3.07), third after F and O. It exhibits typical non-metallic (acidic) properties, forming various oxygen-containing acids, salts and binary compounds, as well as ammonium cation NH 4 and its salts.

In nature - seventeenth by chemical abundance element (ninth among non-metals). A vital element for all organisms.

N 2

Simple substance. It consists of non-polar molecules with a very stable ˚σππ-bond N≡N, this explains the chemical inertness of the element under normal conditions.

A colorless, tasteless and odorless gas that condenses into a colorless liquid (unlike O2).

home component air 78.09% by volume, 75.52 by mass. Nitrogen boils away from liquid air before oxygen does. Slightly soluble in water (15.4 ml/1 l H 2 O at 20 ˚C), the solubility of nitrogen is less than that of oxygen.

At room temperature N2 reacts with fluorine and, to a very small extent, with oxygen:

N 2 + 3F 2 = 2NF 3, N 2 + O 2 ↔ 2NO

The reversible reaction to produce ammonia occurs at a temperature of 200˚C, under pressure up to 350 atm and always in the presence of a catalyst (Fe, F 2 O 3, FeO, in the laboratory with Pt)

N 2 + 3H 2 ↔ 2NH 3 + 92 kJ

According to Le Chatelier's principle, an increase in ammonia yield should occur with increasing pressure and decreasing temperature. However, the reaction rate at low temperatures is very low, so the process is carried out at 450-500 ˚C, achieving a 15% ammonia yield. Unreacted N 2 and H 2 are returned to the reactor and thereby increase the degree of reaction.

Nitrogen is chemically passive in relation to acids and alkalis and does not support combustion.

Receipt V industry– fractional distillation of liquid air or removal of oxygen from air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO when heated. In these cases, nitrogen is obtained, which also contains impurities of noble gases (mainly argon).

In the laboratory, small amounts of chemically pure nitrogen can be obtained by the commutation reaction with moderate heating:

N -3 H 4 N 3 O 2(T) = N 2 0 + 2H 2 O (60-70)

NH 4 Cl(p) + KNO 2 (p) = N 2 0 + KCl + 2H 2 O (100˚C)

Used for ammonia synthesis. Nitric acid and other nitrogen-containing products, as an inert medium for conducting chemical and metallurgical processes and storage of flammable substances.

N.H. 3

Binary compound, the oxidation state of nitrogen is – 3. Colorless gas with a sharp characteristic odor. The molecule has the structure of an incomplete tetrahedron [: N(H) 3 ] (sp 3 hybridization). The presence of a donor pair of electrons in the sp 3 hybrid orbital for nitrogen in the NH 3 molecule determines characteristic reaction addition of a hydrogen cation, thereby forming a cation ammonium NH4. It liquefies under excess pressure at room temperature. In the liquid state, it is associated through hydrogen bonds. Thermally unstable. Highly soluble in water (more than 700 l/1 l H 2 O at 20˚C); the share in a saturated solution is 34% by weight and 99% by volume, pH = 11.8.

Very reactive, prone to addition reactions. Burns in oxygen, reacts with acids. It exhibits reducing (due to N -3) and oxidizing (due to H +1) properties. It is dried only with calcium oxide.

Qualitative reactions – the formation of white “smoke” upon contact with gaseous HCl, blackening of a piece of paper moistened with a solution of Hg 2 (NO3) 2.

An intermediate product in the synthesis of HNO 3 and ammonium salts. Used in the production of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a refrigerant. Poisonous.
Equations of the most important reactions:

2NH 3 (g) ↔ N 2 + 3H 2
NH 3 (g) + H 2 O ↔ NH 3 * H 2 O (p) ↔ NH 4 + + OH —
NH 3 (g) + HCl (g) ↔ NH 4 Cl (g) white “smoke”
4NH 3 + 3O 2 (air) = 2N 2 + 6 H 2 O (combustion)
4NH 3 + 5O 2 = 4NO+ 6 H 2 O (800˚C, cat. Pt/Rh)
2 NH 3 + 3CuO = 3Cu + N 2 + 3 H 2 O (500˚C)
2 NH 3 + 3Mg = Mg 3 N 2 +3 H 2 (600 ˚C)
NH 3 (g) + CO 2 (g) + H 2 O = NH 4 HCO 3 (room temperature, pressure)
Receipt. IN laboratories– displacement of ammonia from ammonium salts when heated with soda lime: Ca(OH) 2 + 2NH 4 Cl = CaCl 2 + 2H 2 O + NH 3
Or boiling an aqueous solution of ammonia and then drying the gas.
In industry Ammonia is produced from nitrogen and hydrogen. Produced by industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.

Ammonia hydrateN.H. 3 * H 2 O. Intermolecular connection. White, in crystal lattice– NH 3 and H 2 O molecules connected by a weak hydrogen bond. Present in an aqueous solution of ammonia, a weak base (dissociation products - NH 4 cation and OH anion). The ammonium cation has a regular tetrahedral structure (sp 3 hybridization). Thermally unstable, completely decomposes when the solution is boiled. Neutralized by strong acids. Shows reducing properties (due to N-3) in a concentrated solution. It undergoes ion exchange and complexation reactions.

Qualitative reaction– formation of white “smoke” upon contact with gaseous HCl. It is used to create a slightly alkaline environment in solution during the precipitation of amphoteric hydroxides.
A 1 M ammonia solution contains mainly NH 3 *H 2 O hydrate and only 0.4% NH 4 OH ions (due to hydrate dissociation); Thus, the ionic “ammonium hydroxide NH 4 OH” is practically not contained in the solution, and there is no such compound in the solid hydrate.
Equations of the most important reactions:
NH 3 H 2 O (conc.) = NH 3 + H 2 O (boiling with NaOH)
NH 3 H 2 O + HCl (diluted) = NH 4 Cl + H 2 O
3(NH 3 H 2 O) (conc.) + CrCl 3 = Cr(OH) 3 ↓ + 3 NH 4 Cl
8(NH 3 H 2 O) (conc.) + 3Br 2(p) = N 2 + 6 NH 4 Br + 8H 2 O (40-50˚C)
2(NH 3 H 2 O) (conc.) + 2KMnO 4 = N 2 + 2MnO 2 ↓ + 4H 2 O + 2KOH
4(NH 3 H 2 O) (conc.) + Ag 2 O = 2OH + 3H 2 O
4(NH 3 H 2 O) (conc.) + Cu(OH) 2 + (OH) 2 + 4H 2 O
6(NH 3 H 2 O) (conc.) + NiCl 2 = Cl 2 + 6H 2 O
A dilute ammonia solution (3-10%) is often called ammonia(the name was invented by alchemists), and the concentrated solution (18.5 - 25%) is an ammonia solution (produced by industry).

Nitrogen oxides

Nitrogen monoxideNO

Non-salt-forming oxide. Colorless gas. Radical, contains a covalent σπ bond (N꞊O), in the solid state a dimer of N 2 O 2 co N-N connection. Extremely thermally stable. Sensitive to air oxygen (turns brown). Slightly soluble in water and does not react with it. Chemically passive towards acids and alkalis. When heated, it reacts with metals and non-metals. a highly reactive mixture of NO and NO 2 (“nitrous gases”). Intermediate product in the synthesis of nitric acid.
Equations of the most important reactions:
2NO + O 2 (g) = 2NO 2 (20˚C)
2NO + C (graphite) = N 2 + CO 2 (400-500˚C)
10NO + 4P(red) = 5N 2 + 2P 2 O 5 (150-200˚C)
2NO + 4Cu = N 2 + 2 Cu 2 O (500-600˚C)
Reactions to mixtures of NO and NO 2:
NO + NO 2 +H 2 O = 2HNO 2 (p)
NO + NO 2 + 2KOH(dil.) = 2KNO 2 + H 2 O
NO + NO 2 + Na 2 CO 3 = 2Na 2 NO 2 + CO 2 (450-500˚C)
Receipt V industry: oxidation of ammonia with oxygen on a catalyst, in laboratories— interaction of dilute nitric acid with reducing agents:
8HNO 3 + 6Hg = 3Hg 2 (NO 3) 2 + 2 NO+ 4 H 2 O
or nitrate reduction:
2NaNO 2 + 2H 2 SO 4 + 2NaI = 2 NO + I 2 ↓ + 2 H 2 O + 2Na 2 SO 4

Nitrogen dioxideNO 2

Acid oxide, conditionally corresponds to two acids - HNO 2 and HNO 3 (acid for N 4 does not exist). Brown gas, at room temperature a monomer NO 2, in the cold a liquid colorless dimer N 2 O 4 (dianitrogen tetroxide). Reacts completely with water and alkalis. A very strong oxidizing agent that causes corrosion of metals. It is used for the synthesis of nitric acid and anhydrous nitrates, as a rocket fuel oxidizer, an oil purifier from sulfur, and a catalyst for the oxidation of organic compounds. Poisonous.
Equation of the most important reactions:
2NO 2 ↔ 2NO + O 2
4NO 2 (l) + H 2 O = 2HNO 3 + N 2 O 3 (syn.) (in the cold)
3 NO 2 + H 2 O = 3HNO 3 + NO
2NO 2 + 2NaOH (diluted) = NaNO 2 + NaNO 3 + H 2 O
4NO 2 + O 2 + 2 H 2 O = 4 HNO 3
4NO 2 + O 2 + KOH = KNO 3 + 2 H 2 O
2NO 2 + 7H 2 = 2NH 3 + 4 H 2 O (cat. Pt, Ni)
NO 2 + 2HI(p) = NO + I 2 ↓ + H 2 O
NO 2 + H 2 O + SO 2 = H 2 SO 4 + NO (50-60˚C)
NO 2 + K = KNO 2
6NO 2 + Bi(NO 3) 3 + 3NO (70-110˚C)
Receipt: V industry - oxidation of NO by atmospheric oxygen, in laboratories– interaction of concentrated nitric acid with reducing agents:
6HNO 3 (conc., hor.) + S = H 2 SO 4 + 6NO 2 + 2H 2 O
5HNO 3 (conc., hor.) + P (red) = H 3 PO 4 + 5NO 2 + H 2 O
2HNO 3 (conc., hor.) + SO 2 = H 2 SO 4 + 2 NO 2

Dianitrogen oxideN 2 O

A colorless gas with a pleasant odor (“laughing gas”), N꞊N꞊О, formal oxidation state of nitrogen +1, poorly soluble in water. Supports combustion of graphite and magnesium:

2N 2 O + C = CO 2 + 2N 2 (450˚C)
N 2 O + Mg = N 2 + MgO (500˚C)
Obtained by thermal decomposition of ammonium nitrate:
NH 4 NO 3 = N 2 O + 2 H 2 O (195-245˚C)
used in medicine as an anesthetic.

Dianitrogen trioxideN 2 O 3

At low temperatures – blue liquid, ON꞊NO 2, formal oxidation state of nitrogen +3. At 20 ˚C, it decomposes 90% into a mixture of colorless NO and brown NO 2 (“nitrous gases”, industrial smoke – “fox tail”). N 2 O 3 is an acidic oxide, in the cold with water it forms HNO 2, when heated it reacts differently:
3N 2 O 3 + H 2 O = 2HNO 3 + 4NO
With alkalis it gives salts HNO 2, for example NaNO 2.
Obtained by reacting NO with O 2 (4NO + 3O 2 = 2N 2 O 3) or with NO 2 (NO 2 + NO = N 2 O 3)
with strong cooling. “Nitrous gases” are also environmentally dangerous and act as catalysts for the destruction of the ozone layer of the atmosphere.

Dianitrogen pentoxide N 2 O 5

Colorless, solid substance, O 2 N – O – NO 2, nitrogen oxidation state is +5. At room temperature it decomposes into NO 2 and O 2 in 10 hours. Reacts with water and alkalis as an acid oxide:
N2O5 + H2O = 2HNO3
N 2 O 5 + 2NaOH = 2NaNO 3 + H 2
Prepared by dehydration of fuming nitric acid:
2HNO3 + P2O5 = N2O5 + 2HPO3
or oxidation of NO 2 with ozone at -78˚C:
2NO 2 + O 3 = N 2 O 5 + O 2

Nitrites and nitrates

Potassium nitriteKNO 2 . White, hygroscopic. Melts without decomposition. Stable in dry air. Very soluble in water (forming a colorless solution), hydrolyzes at the anion. A typical oxidizing and reducing agent in an acidic environment, it reacts very slowly in an alkaline environment. Enters into ion exchange reactions. Qualitative reactions on the NO 2 ion - discoloration of the violet MnO 4 solution and the appearance of a black precipitate when adding I ions. It is used in the production of dyes, as an analytical reagent for amino acids and iodides, and a component of photographic reagents.
equation of the most important reactions:
2KNO 2 (t) + 2HNO 3 (conc.) = NO 2 + NO + H 2 O + 2KNO 3
2KNO 2 (dil.)+ O 2 (e.g.) → 2KNO 3 (60-80 ˚C)
KNO 2 + H 2 O + Br 2 = KNO 3 + 2HBr
5NO 2 - + 6H + + 2MnO 4 - (viol.) = 5NO 3 - + 2Mn 2+ (bts.) + 3H 2 O
3 NO 2 - + 8H + + CrO 7 2- = 3NO 3 - + 2Cr 3+ + 4H 2 O
NO 2 - (saturated) + NH 4 + (saturated) = N 2 + 2H 2 O
2NO 2 - + 4H + + 2I - (bts.) = 2NO + I 2 (black) ↓ = 2H 2 O
NO 2 - (diluted) + Ag + = AgNO 2 (light yellow)↓
Receipt Vindustry– reduction of potassium nitrate in the processes:
KNO3 + Pb = KNO 2+ PbO (350-400˚C)
KNO 3 (conc.) + Pb (sponge) + H 2 O = KNO 2+ Pb(OH) 2 ↓
3 KNO3 + CaO + SO2 = 2 KNO 2+ CaSO 4 (300 ˚C)

H itrate potassium KNO 3
Technical name potash, or Indian salt , saltpeter. White, melts without decomposition and decomposes upon further heating. Stable in air. Highly soluble in water (with high endo-effect, = -36 kJ), no hydrolysis. A strong oxidizing agent during fusion (due to the release of atomic oxygen). In solution it is reduced only by atomic hydrogen (in an acidic environment to KNO 2, in an alkaline environment to NH 3). Used in glass production as a preservative food products, a component of pyrotechnic mixtures and mineral fertilizers.

2KNO 3 = 2KNO 2 + O 2 (400-500 ˚C)

KNO 3 + 2H 0 (Zn, dil. HCl) = KNO 2 + H 2 O

KNO 3 + 8H 0 (Al, conc. KOH) = NH 3 + 2H 2 O + KOH (80 ˚C)

KNO 3 + NH 4 Cl = N 2 O + 2H 2 O + KCl (230-300 ˚C)

2 KNO 3 + 3C (graphite) + S = N 2 + 3CO 2 + K 2 S (combustion)

KNO 3 + Pb = KNO 2 + PbO (350 - 400 ˚C)

KNO 3 + 2KOH + MnO 2 = K 2 MnO 4 + KNO 2 + H 2 O (350 - 400 ˚C)

Receipt: in industry
4KOH (hor.) + 4NO 2 + O 2 = 4KNO 3 + 2H 2 O

and in the laboratory:
KCl + AgNO 3 = KNO 3 + AgCl↓

Individual characteristics organic matter is its IR spectrum.

It should be noted that even a novice researcher can often only draw a conclusion from IR and UV spectroscopy that a substance belongs to a particular class of chemical compounds, without carrying out painstaking research. chemical reactions. The problem is solved extremely simply in most cases using PMR spectroscopy.

Identification of a compound is carried out by establishing the identity of the constants (Tm, Tbp, Rf,nD etc.) both the compound being determined and its derivatives with constants of known substances given in the table of derivatives for identification. The study is carried out in the following sequence.

The physical properties of the compound are studied: state of aggregation, color, smell, boiling and melting points, solubility and relationship to calcination. Using these data, it is sometimes possible to immediately determine the class of compounds to which the analyzed substance belongs, significantly reducing the number of subsequent operations.

The qualitative composition is determined, i.e., samples are taken for the presence of carbon, hydrogen, nitrogen, halogens, and sulfur. In the absence of one or another of the listed elements, it is possible not to make qualitative reactions to the functional groups containing them. (Knowledge of the quantitative elemental composition of a substance is of great benefit in identifying a compound. However, quantitative elemental analysis can only be performed in a specially equipped laboratory or on automatic instruments - C-, H-, N-analyzers.)

Individual functional groups are discovered and the class of the analyte is established.

The substance is converted into one or two derivatives that are most characteristic of a given class, using the constants of which a conclusion is drawn about the exact structure of the substance (identification).

Confirm the structure of the compound by taking or obtaining UV, IR and PMR spectra of the substance from the teacher.

1.Preliminary study of the substance

The study begins with a description of the external properties of the substance: state of aggregation, color, smell. If the substance is solid, then write down what it is - crystalline (needles, plates, prisms, etc.), microcrystalline, amorphous. Pay attention to whether this substance is homogeneous or not. Record its color. Colored ones include quinones, some α-diketones, azo-, nitroso-, nitro derivatives, some polyhalogen derivatives, compounds with a large number conjugated bonds. It is noted whether the color is retained during recrystallization and distillation, i.e., whether it is inherent in the substance or due to impurities. For crystalline substances, the melting point is determined; for liquid substances, the boiling point and refractive index of light are determined.

Write down the constants of the substance; color and smell of a pure substance. Many organic compounds have a specific odor, by which, with skill, one can determine which class they belong to (ethers, phenols, nitro compounds, amines, etc.).

Calcination. Place 0.1 ml (0.1 g) of the substance on the crucible lid (solid on the tip of the spatula). Carefully introduce it into the upper or side part of the colorless flame of the burner, gradually moving the lid into the hotter part of the flame. Carefully observe the changes occurring in the substance. Record the nature of melting (whether the substance decomposes) and combustion (fast, with a flash, slow), the color of the flame, and the smell. If a substance burns with a dim flame (almost blue), this indicates the presence of oxygen-containing functional groups in it. A yellow glowing (smoking) flame is characteristic of carbon-rich compounds (aromatic and acetylene hydrocarbons).

Determination of solubility. Based on the solubility of a substance in various solvents, one can draw a conclusion about the presence of certain functional groups in the substance. In addition, determining solubility allows you to select a suitable solvent for recrystallization of a substance (like dissolves in like). It is advisable to study solubility in the following solvents: water, 6% solutions of sodium hydroxide, sodium bicarbonate, of hydrochloric acid; concentrated sulfuric acid, ethyl alcohol, benzene, petroleum ether, acetic acid. Add one drop of liquid or 0.01 g of solid compound and drop by drop -0.2 ml (10 drops) of solvent into the test tube. After each added portion of the solvent, the test tube is shaken. If a compound is completely soluble, it is recorded as soluble. If the substance is poorly soluble or does not dissolve at room temperature, heat the mixture to a boil. In case of poor solubility in inorganic solvents, the undissolved substance is separated, and the solution is neutralized and observed whether the original compound is released from it. The turbidity of the neutralized filtrate indicates the properties of the substances: acidic if the solvent was alkali or soda; basic - acidic solvent. When adding a substance to a hydrogen carbonate solution, you need to pay attention to whether carbon monoxide (IV) is released.

2. Qualitative analysis

Nitrogen, sulfur, halogen can be detected in one portion of the substance by fusing it with sodium (if the substance is liquid, first make sure that it is not an acid, otherwise an explosion is possible):

CnHmHalNS → NaHal + NaCN + Na2S

Having dissolved the alloy, qualitative reactions are carried out on the following ions: Hal-, S2-, CN-. ~0.1 g of the substance is placed in a test tube. Secure the test tube in a rack at an angle in a fume hood. Add a small piece (about a quarter of a pea) of purified sodium to it. Carefully heat to a dark red heat and quickly lower the test tube into a glass with 5 ml of distilled water ~Be careful, wear glasses! ~ The solution is filtered from glass fragments and poured into several test tubes into separate portions of 1-1.5 ml. Each portion is used to carry out a qualitative reaction (the filtrate should be colorless). Belstein tests are done with silver nitrate for halogens, with lead acetate for sulfur and nitrogen.↓

Qualitative reactions to halogens

Belstein test. Halogens can be detected in the substance under study without fusing it with sodium. The end of the copper wire is bent into a small loop and calcined in the flame of the burner until the green color of the flame disappears. Allow the wire to cool, immerse it in the test substance and heat it again in the flame of the burner. A green flame indicates the presence of halogens. This reaction has exceptionally high sensitivity (impurities can also give a positive reaction! Therefore, its positive result should always be double-checked by reaction with silver nitrate):

AgNO3 + NaHal → AgHal + NaNO3

The filtrate obtained after decomposition of the test substance by fusion with sodium is acidified with nitric acid to an acidic reaction and added water solution silver nitrate. A cheesy precipitate of silver halide of white (chlorine), yellowish (bromine) and bright yellow (iodine) color appears.

Qualitative reaction to nitrogen

FeSO4 + 2NaCN → Fe(CN)2 + Na2S04

Fe(CN)2 + 4NaCN → Na4Fe(CN)6

3Nа4Fe(СN)6 + 2Fe2(SO4)З → Fe4З + 6Na2SO4

A crystal of iron sulfate or 2 drops of a freshly prepared solution is added to the filtrate. Boil for one minute. Add a drop of Fe3+ salt solution. Acidify with hydrochloric acid (5-6 drops). If nitrogen is present in the test substance, a Prussian blue precipitate appears or a bright blue color appears.

Qualitative reaction to sulfur

Na2S + 2HCl → H2S + 2NaCl (a)

Na2S + Pb(OCOC3)2 → PbS↓ + 2СН3СООНа (b)

Na2S + Na2 → Na4 (c)

To detect the S2- ion, one portion of the filtrate is acidified with hydrochloric acid. The characteristic smell of hydrogen sulfide will indicate the presence of divalent sulfur (a). In another test tube, the filtrate is acidified with acetic acid and a solution of lead acetate is added. In the presence of S2-, a black precipitate PbS is formed. In the case of a small amount of sulfur, instead of precipitation, the solution only turns brown (b). To the third portion of the filtrate add a few drops of a diluted solution of sodium nitroprusside. The appearance of a blue-violet color of the thionitro complex indicates the presence of sulfur (c).

3. Opening functional groups

Based on the research results physical properties And quality composition compounds, determine the approximate possible class of the analyte. Qualitative reactions are then made for the putative functional groups. Let's say it is established: the substance is liquid, colorless, does not contain nitrogen, halogens and sulfur, dissolves well in water, has a neutral reaction, boils at 78 ° C. Presumably, such a substance could be an alcohol, aldehyde, or ketone. To clarify, qualitative reactions are made only for alcohol, aldehyde and ketone groups. Small samples (0.1-0.15 g) of the substance should be taken, retaining the bulk for obtaining derivatives and (part as a reserve) for final specific reactions to a given individual substance.

Before carrying out any reaction with the analyte, it is advisable to conduct an experiment with a known compound of this class. And only after mastering the technique of performing the operation and making sure good quality reagents, you should proceed to samples with the analyte.

Multiple connection

Reaction with bromine.

The vast majority of compounds containing a multiple bond (double, triple, or combinations thereof, with the exception of aromatic systems), bromine is easily added:

The reaction is usually carried out in acetic acid or carbon tetrachloride. To a solution of 0.1 g or 0.1 ml of the substance in 2-3 ml of glacial acetic acid, placed in a small test tube, add dropwise, shaking, a 1% solution of bromine in glacial acetic acid. If there is a multiple bond in the substance, the solution instantly becomes colorless.

In some cases, compounds containing hydrogen, which is easily replaced by bromine (aniline, phenol, ketones, some tertiary hydrocarbons), also discolor the bromine solution. However, this releases hydrogen bromide, which is easily determined using wet litmus paper or Congo:

2.1.1. Qualitative reactions to the sulfide anion S 2-. Of the sulfides, only sulfides of alkali metals and ammonium are soluble. Insoluble sulfides have a specific color, by which one or another sulfide can be identified.
Color:
MnS - flesh (pink).
ZnS - white.
PbS - black.
Ag 2 S - black.
CdS - lemon yellow.
SnS - chocolate.
HgS (metacinnabar) - black.
HgS (cinnabar) - red.
Sb 2 S 3 - orange.
Bi 2 S 3 - black.
Some sulfides, when interacting with non-oxidizing acids, form a toxic gas, hydrogen sulfide H 2 S, with an unpleasant odor (rotten eggs):
Na 2 S + 2HBr = 2NaBr + H 2 S
S 2- + 2H + = H 2 S

And some are resistant to dilute solutions of HCl, HBr, HI, H 2 SO 4, HCOOH, CH 3 COOH - for example CuS, Cu 2 S, Ag 2 S, HgS, PbS, CdS, Sb 2 S 3, SnS and some others . But they are transferred into a conc. solution. nitric acid when boiling (Sb 2 S 3 and HgS dissolve the hardest, and the latter will dissolve much faster in aqua regia):
CuS + 8HNO 3 =t= CuSO 4 + 8NO 2 + 4H 2 O

The sulfide anion can also be identified by adding a sulfide solution to bromine water:
S 2- + Br 2 = S↓ + 2Br -
The resulting sulfur precipitates.

2.1.2. Qualitative reaction to the sulfate anion SO 4 2-. The sulfate anion is usually precipitated with a lead or barium cation:
Pb 2+ + SO 4 2- = PbSO 4 ↓
Lead sulfate precipitate is white.

2.1.3. Qualitative reaction to the silicate anion SiO 3 2-. The silicate anion easily precipitates from solution in the form of a glassy mass when strong acids are added:
SiO 3 2- + 2H + = H 2 SiO 3 ↓ (SiO 2 * nH 2 O)

2.1.4. Qualitative reactions to the chloride anion Cl -, bromide anion Br -, iodide anion I - see the paragraph “qualitative reactions to the silver cation Ag +”.

2.1.5. Qualitative reaction to the sulfite anion SO 3 2-. When strong acids are added to a solution, sulfur dioxide SO2 is formed - a gas with a pungent odor (the smell of a lit match):
SO 3 2- + 2H + = SO 2 + H 2 O

2.1.6. Qualitative reaction to the carbonate anion CO 3 2-. When strong acids are added to a carbonate solution, carbon dioxide CO 2 is formed, which extinguishes the burning splinter:
CO 3 2- + 2H + = CO 2 + H 2 O

2.1.7. Qualitative reaction to the thiosulfate anion S 2 O 3 2-. When a solution of sulfuric or hydrochloric acid is added to a solution of thiosulfate, sulfur dioxide SO2 is formed and elemental sulfur S precipitates:
S 2 O 3 2- + 2H + = S↓ + SO 2 + H2O

2.1.8. Qualitative reaction to the chromate anion CrO 4 2-. When a solution of barium salts is added to a chromate solution, a yellow precipitate of barium chromate BaCrO 4 precipitates, decomposing in a strongly acidic environment:
Ba 2+ + CrO 4 2- = BaCrO 4 ↓
Chromate solutions are colored yellow. When the solution is acidified, the color changes to orange, corresponding to the dichromate anion Cr 2 O 7 2-:
2CrO 4 2- + 2H + = Cr 2 O 7 2- + H 2 O
In addition, chromates are oxidizing agents in alkaline and neutral environments (oxidizing abilities are worse than those of dichromates):
S 2- + CrO 4 2- + H 2 O = S + Cr(OH) 3 + OH -

2.1.9. Qualitative reaction to the dichromate anion Cr 2 O 7 2-. When a silver salt solution is added to a dichromate solution, an orange precipitate Ag 2 Cr 2 O 7 is formed:
2Ag + + Cr 2 O 7 2- = Ag 2 Cr 2 O 7 ↓
Solutions of dichromates are orange. When the solution is alkalized, the color changes to yellow, corresponding to the chromate anion CrO 4 2-:
Cr 2 O 7 2- + 2OH - = 2CrO 4 2- + H 2 O
In addition, dichromates are strong oxidizing agents in an acidic environment. When any reducing agent is added to an acidified dichromate solution, the color of the solution will change from orange to green, corresponding to the chromium (III) cation Cr 3+ (bromide anion as a reducing agent):
6Br - + Cr 2 O 7 2- + 14H + = 3Br 2 + 2Cr 3+ + 7H 2 O
A spectacular qualitative reaction to hexavalent chromium is a dark blue coloration of the solution when conc. hydrogen peroxide in ether. Chromium peroxide of the composition CrO 5 is formed.

2.2.0. Qualitative reaction to the permanganate anion MnO 4 -. The permanganate anion “gives out” the dark purple color of the solution. In addition, permanganates are the strongest oxidizing agents; in an acidic environment they are reduced to Mn 2+ (the purple color disappears), in a neutral environment - to Mn +4 (the color disappears, a brown precipitate of manganese dioxide MnO 2 precipitates) and in an alkaline environment - to MnO 4 2- (the color of the solution changes to dark green):
5SO 3 2- + 2MnO 4 - + 6H + = 5SO 4 2- + 2Mn 2+ + 3H 2 O
3SO 3 2- + 2MnO 4 - + H 2 O = 3SO 4 2- + 2MnO 2 ↓ + 2OH -
SO 3 2- + 2MnO 4 - + 2OH - = SO 4 2- + 2MnO 4 2- + H 2 O

2.2.1. Qualitative reaction to the manganate anion MnO 4 2-. When the manganate solution is acidified, the dark green color changes to dark purple, corresponding to the permanganate anion MnO 4 -:
3K 2 MnO 4 (r.) + 4HCl (dil.) = MnO 2 ↓ + 2KMnO 4 + 4KCl + 2H 2 O

2.2.2. Qualitative reaction to the phosphate anion PO 4 3-. When a silver salt solution is added to a phosphate solution, a yellowish precipitate of silver (I) phosphate Ag 3 PO 4 precipitates:
3Ag + + PO 4 3- = Ag 3 PO 4 ↓
The reaction to the dihydrogen phosphate anion H 2 PO 4 - is similar.

2.2.3. Qualitative reaction to the ferrate anion FeO 4 2-. Precipitation of red barium ferrate from a solution (the reaction is carried out in an alkali environment):
Ba 2+ + FeO 4 2- =OH - = BaFeO 4 ↓
Ferrates are the strongest oxidizing agents (stronger than permanganates). Stable in alkaline environment, unstable in acidic environment:
4FeO 4 2- + 20H + = 4Fe 3+ + 3O 2 + 10H 2 O

2.2.4. Qualitative reaction to the nitrate anion NO 3 -. Nitrates in solution do not exhibit oxidizing properties. But when the solution is acidified, they can oxidize, for example, copper (the solution is usually acidified with diluted H 2 SO 4):
3Cu + 2NO 3 - + 8H + = 3Cu 2+ + 2NO + 4H 2 O

2.2.5. Qualitative reaction to hexacyanoferrate (II) and (III) ions 4- and 3-. When adding solutions containing Fe 2+, a dark blue precipitate is formed (Turnboole blue, Prussian blue):
K 3 + FeCl 2 = KFe + 2KCl (in this case, the precipitate consists of a mixture of KFe(II), KFe(III), Fe 3 2, Fe 4 3).

2.2.6. Qualitative reaction to the arsenate anion AsO 4 3-. Formation of water-insoluble silver (I) arsenate Ag 3 AsO 4, which has a “cafe au lait” color:
3Ag + + AsO 4 3- = Ag 3 AsO 4 ↓
Here are the main qualitative reactions to anions. Next we will look at qualitative reactions to simple and complex substances.

3. Qualitative reactions to simple and complex substances. Some simple and complex substances, like ions, are detected by qualitative reactions. Below I will describe qualitative reactions to some substances.

3.1.1. Qualitative reaction to hydrogen H2. A barking pop when you bring a burning splinter to a source of hydrogen.

3.1.2. Qualitative reaction to nitrogen N2. Extinguishing a burning splinter in a nitrogen atmosphere. When Ca(OH) 2 is passed into the solution, no precipitate forms.

3.1.3. Qualitative reaction to oxygen O 2. The bright ignition of a smoldering splinter in an oxygen atmosphere.

3.1.4. Qualitative reaction to ozone O 3. The interaction of ozone with a solution of iodides with the precipitation of crystalline iodine I 2 into a precipitate:
2KI + O 3 + H 2 O = 2KOH + I 2 ↓ + O 2
Unlike ozone, oxygen in this reaction Not enters.

3.1.5. Qualitative reaction to chlorine Cl 2.Chlorine is a yellow-green gas with a very unpleasant odor. interaction of a lack of chlorine with solutions of iodides, elemental iodine I 2 precipitates:
2KI + Cl 2 = 2KCl + I 2 ↓
Excess chlorine will lead to oxidation of the resulting iodine:
I 2 + 5Cl 2 + 6H 2 O = 2HIO 3 + 10HCl

3.1.6. Qualitative reactions to ammonia NH 3. Note: these reactions are not given in the school course. However, these are the most reliable qualitative reactions to ammonia.
Blackening of a piece of paper soaked in a solution of mercury salt (I) Hg 2 +:
Hg 2 Cl 2 + 2NH 3 = Hg(NH 2)Cl + Hg + NH 4 Cl
The paper turns black due to the release of fine mercury.

Interaction of ammonia with an alkaline solution of potassium tetraiodomercurate (II) K 2 (Nessler reagent) :
2K2 + NH3 + 3KOH = I · H 2 O↓ + 7KI + 2H 2 O
Complex I · H 2 O is brown in color (rust color) and precipitates.
The last two reactions are the most reliable for ammonia.

Reaction of ammonia with hydrogen chloride (“smoke” without fire):
NH 3 + HCl = NH 4 Cl

3.1.7. Qualitative reaction to phosgene (carbon chloride, carbonyl chloride) COCl 2. Emission of white “smoke” from a piece of paper soaked in ammonia solution:
COCl 2 + 4NH 3 = (NH 2) 2 CO + 2NH 4 Cl

3.1.8. Qualitative reaction to carbon monoxide (carbon monoxide) CO. Cloudiness of the solution when passing through carbon monoxide into a solution of palladium (II) chloride:
PdCl 2 + CO + H 2 O = CO 2 + 2HCl + Pd↓

3.1.9. Qualitative reaction to carbon dioxide (carbon dioxide) CO 2. Extinguishing a smoldering splinter in the atmosphere carbon dioxide.
Passing carbon dioxide into a solution of slaked lime Ca(OH) 2:
Ca(OH) 2 + CO 2 = CaCO 3 ↓ + H 2 O
Further passing will lead to the dissolution of the precipitate:
CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 2

3.2.1. Qualitative reaction to nitric oxide (II) NO. Nitrogen oxide (II) is very sensitive to atmospheric oxygen, therefore it turns brown in air, oxidizing to nitrogen oxide (IV) NO 2:
2NO + O 2 = 2NO 2

2.1. Qualitative reactions to the sulfide anion S 2 -. Of the sulfides, only sulfides of alkali metals and ammonium are soluble. Insoluble sulfides have a specific color, by which one or another sulfide can be identified.
Color:
MnS - flesh (pink).
ZnS - white.
PbS - black.
Ag 2 S - black.
CdS - lemon yellow.
SnS - chocolate.
HgS (metacinnabar) - black.
HgS (cinnabar) - red.
Sb 2 S 3 - orange.
Bi 2 S 3 - black.
Some sulfides, when interacting with non-oxidizing acids, form a toxic gas, hydrogen sulfide H 2 S, with an unpleasant odor (rotten eggs):
Na 2 S + 2HBr = 2NaBr + H 2 S
S 2- + 2H + = H 2 S

And some are resistant to dilute solutions of HCl, HBr, HI, H 2 SO 4, HCOOH, CH 3 COOH - for example CuS, Cu 2 S, Ag 2 S, HgS, PbS, CdS, Sb 2 S 3, SnS and some others . But they are transferred into a conc. solution. nitric acid when boiling (Sb 2 S 3 and HgS dissolve the hardest, and the latter will dissolve much faster in aqua regia):
CuS + 8HNO 3 =t= CuSO 4 + 8NO 2 + 4H 2 O

The sulfide anion can also be identified by adding a sulfide solution to bromine water:
S 2- + Br 2 = S↓ + 2Br -
The resulting sulfur precipitates.

2.2. Qualitative reaction to the sulfate anion SO 4 2- . The sulfate anion is usually precipitated with a lead or barium cation:
Pb 2+ + SO 4 2- = PbSO 4 ↓

Ba 2+ + SO 4 2- = BaSO 4 ↓
Precipitates of lead sulfate and barium sulfate are white.

2.3. Qualitative reaction to the silicate anion SiO 3 2-. The silicate anion easily precipitates from solution in the form of a glassy mass when strong acids are added:
SiO 3 2- + 2H + = H 2 SiO 3 ↓ (SiO 2 * nH 2 O)

2.4. Qualitative reactions to the chloride anion Cl -, bromide anion Br -, iodide anion I - see paragraph “qualitative reactions to silver cation Ag +”

2.5. Qualitative reaction to the sulfite anion SO 3 2-. When strong acids are added to a solution, sulfur dioxide SO2 is formed - a gas with a pungent odor (the smell of a lit match):
SO 3 2- + 2H + = SO 2 + H 2 O

2.6. Qualitative reaction to the carbonate anion CO 3 2-. When strong acids are added to a carbonate solution, carbon dioxide CO 2 is formed, which does not support combustion and causes clouding of lime water:
CO 3 2- + 2H + = CO 2 + H 2 O

Ca(OH) 2 + CO 2 = CaCO 3 ↓ + H 2 O (with an excess of CO 2, the precipitate dissolves to form bicarbonate CaCO 3 + CO 2 + H 2 O = Ca (HCO 3) 2

2.7. Qualitative reaction to the thiosulfate anion S 2 O 3 2-. When a solution of sulfuric or hydrochloric acid is added to a solution of thiosulfate, sulfur dioxide SO2 is formed and elemental sulfur S precipitates:
S 2 O 3 2- + 2H + = S↓ + SO 2 + H2O

2.8. Qualitative reaction to the chromate anion CrO 4 2-. When a solution of barium salts is added to a chromate solution, a yellow precipitate of barium chromate BaCrO 4 precipitates, decomposing in a strongly acidic environment:
Ba 2+ + CrO 4 2- = BaCrO 4 ↓
Chromate solutions are colored yellow. When the solution is acidified, the color changes to orange, corresponding to the dichromate anion Cr 2 O 7 2-:
2CrO 4 2- + 2H + = Cr 2 O 7 2- + H 2 O
In addition, chromates are oxidizing agents in alkaline and neutral environments (oxidizing abilities are worse than those of dichromates):
S 2- + CrO 4 2- + H 2 O = S + Cr(OH) 3 + OH -

2.9. Qualitative reaction to the dichromate anion Cr 2 O 7 2-. When a silver salt solution is added to a dichromate solution, an orange precipitate Ag 2 Cr 2 O 7 is formed:
2Ag + + Cr 2 O 7 2- = Ag 2 Cr 2 O 7 ↓
Solutions of dichromates are orange. When the solution is alkalized, the color changes to yellow, corresponding to the chromate anion CrO 4 2-:
Cr 2 O 7 2- + 2OH - = 2CrO 4 2- + H 2 O
In addition, dichromates are strong oxidizing agents in an acidic environment. When any reducing agent is added to an acidified dichromate solution, the color of the solution will change from orange to green, corresponding to the chromium (III) cation Cr 3+ (bromide anion as a reducing agent):
6Br - + Cr 2 O 7 2- + 14H + = 3Br 2 + 2Cr 3+ + 7H 2 O
A spectacular qualitative reaction to hexavalent chromium is a dark blue coloration of the solution when conc. hydrogen peroxide in ether. Chromium peroxide of the composition CrO 5 is formed.

2.10. Qualitative reaction to the permanganate anion MnO 4 -. The permanganate anion “gives out” the dark purple color of the solution. In addition, permanganates are the strongest oxidizing agents; in an acidic environment they are reduced to Mn 2+ (the purple color disappears), in a neutral environment - to Mn +4 (the color disappears, a brown precipitate of manganese dioxide MnO 2 precipitates) and in an alkaline environment - to MnO 4 2- (the color of the solution changes to dark green):
5SO 3 2- + 2MnO 4 - + 6H + = 5SO 4 2- + 2Mn 2+ + 3H 2 O
3SO 3 2- + 2MnO 4 - + H 2 O = 3SO 4 2- + 2MnO 2 ↓ + 2OH -
SO 3 2- + 2MnO 4 - + 2OH - = SO 4 2- + 2MnO 4 2- + H 2 O

2.11. Qualitative reaction to the manganate anion MnO 4 2-. When the manganate solution is acidified, the dark green color changes to dark purple, corresponding to the permanganate anion MnO 4 -:
3K 2 MnO 4 (r.) + 4HCl (dil.) = MnO 2 ↓ + 2KMnO 4 + 4KCl + 2H 2 O

2.12. Qualitative reaction to the phosphate anion PO 4 3-. When a silver salt solution is added to a phosphate solution, a yellowish precipitate of silver (I) phosphate Ag 3 PO 4 precipitates:
3Ag + + PO 4 3- = Ag 3 PO 4 ↓
The reaction to the dihydrogen phosphate anion H 2 PO 4 - is similar.

2.13. Qualitative reaction to the ferrate anion FeO 4 2-. Precipitation of red barium ferrate from a solution (the reaction is carried out in an alkali environment):
Ba 2+ + FeO 4 2- =OH - = BaFeO 4 ↓
Ferrates are the strongest oxidizing agents (stronger than permanganates). Stable in alkaline environment, unstable in acidic environment:
4FeO 4 2- + 20H + = 4Fe 3+ + 3O 2 + 10H 2 O

2.14. Qualitative reaction to the nitrate anion NO 3 -. Nitrates in solution do not exhibit oxidizing properties. But when the solution is acidified, they can oxidize, for example, copper (the solution is usually acidified with diluted H 2 SO 4):
3Cu + 2NO 3 - + 8H + = 3Cu 2+ + 2NO + 4H 2 O

2.15. Qualitative reaction to hexacyanoferrate (II) and (III) ions 4- and 3-. When adding solutions containing Fe 2+, a dark blue precipitate is formed (Turnboole blue, Prussian blue):
K 3 + FeCl 2 = KFe + 2KCl (in this case, the precipitate consists of a mixture of KFe(II), KFe(III), Fe 3 2, Fe 4 3).

2.1.6. Qualitative reaction to the arsenate anion AsO 4 3-. Formation of water-insoluble silver (I) arsenate Ag 3 AsO 4, which has a “cafe au lait” color:
3Ag + + AsO 4 3- = Ag 3 AsO 4 ↓

3. Qualitative reactions to simple and complex substances. Some simple and complex substances, like ions, are detected by qualitative reactions.

3.1. Qualitative reaction to hydrogen H2. A characteristic pop when bringing a burning splinter to a source of hydrogen.

3. 2. Qualitative reaction to nitrogen N2. Does not support combustion. When lime water is passed through a solution, no precipitate forms.

3. 3. Qualitative reaction to oxygen O 2. The bright flash of a smoldering splinter in an oxygen atmosphere.

3. 4. Qualitative reaction to ozone O 3. The interaction of ozone with a solution of iodides with the precipitation of crystalline iodine I 2 into a precipitate:
2KI + O 3 + H 2 O = 2KOH + I 2 ↓ + O 2
Oxygen to this reaction Not enters.

3. 5. Qualitative reaction to chlorine Cl 2. Chlorine is a yellow-green gas with a pungent odor. At interaction of a lack of chlorine with solutions of iodides, iodine I 2 precipitates:
2KI + Cl 2 = 2KCl + I 2 ↓
Excess chlorine will lead to oxidation of the resulting iodine:
I 2 + 5Cl 2 + 6H 2 O = 2HIO 3 + 10HCl

3. 6. Qualitative reactions to ammonia NH 3. Note: these reactions are not given in the school course. However, these are the most reliable qualitative reactions to ammonia.
Blackening of a piece of paper soaked in a solution of mercury salt (I) Hg 2 +:
Hg 2 Cl 2 + 2NH 3 = Hg(NH 2)Cl + Hg + NH 4 Cl
The paper turns black due to the release of fine mercury.

Interaction of ammonia with an alkaline solution of potassium tetraiodomercurate (II) K 2 (Nessler reagent) :
2K2 + NH3 + 3KOH = I · H 2 O↓ + 7KI + 2H 2 O
Complex I · H 2 O is brown in color (rust color) and precipitates.
The last two reactions are the most reliable for ammonia.

Reaction of ammonia with hydrogen chloride (“smoke” without fire):
NH 3 + HCl = NH 4 Cl

3. 7. Qualitative reaction to carbon monoxide (carbon monoxide) CO. Cloudiness of the solution when carbon monoxide is passed into a solution of palladium (II) chloride:
PdCl 2 + CO + H 2 O = CO 2 + 2HCl + Pd↓

3. 8. Qualitative reaction to carbon dioxide (carbon dioxide) CO 2. Extinguishing a smoldering splinter in an atmosphere of carbon dioxide.
Passing carbon dioxide into a solution of slaked lime Ca(OH) 2:
Ca(OH) 2 + CO 2 = CaCO 3 ↓ + H 2 O
Further passing will lead to the dissolution of the precipitate:
CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 2

3.9. Qualitative reaction to nitric oxide (II) NO. Nitrogen oxide (II) is very sensitive to atmospheric oxygen, therefore it turns brown in air, oxidizing to nitrogen oxide (IV) NO 2:
2NO + O 2 = 2NO 2