HNO 3 is a strong acid. Its salts - nitrates-- obtained by the action of HNO 3 on metals, oxides, hydroxides or carbonates. All nitrates are highly soluble in water.
Salts of nitric acid - nitrates - decompose irreversibly when heated, the decomposition products are determined by the cation:
NH 4 NO 3 = N 2 O + 2H 2 O
Nitrates in aqueous solutions exhibit practically no oxidizing properties, but when high temperature in the solid state, nitrates are strong oxidizing agents, for example:
Fe + 3KNO 3 + 2KOH = K 2 FeO 4 + 3KNO 2 + H 2 O - when fusing solids.
Zinc And aluminum in an alkaline solution, nitrates are reduced to NH 3:
Nitric acid salts -- nitrates-- widely used as fertilizers. Moreover, almost all nitrates are highly soluble in water, so there are extremely few of them in nature in the form of minerals; the exception is Chilean (sodium) saltpeter and Indian saltpeter ( potassium nitrate). Most nitrates are obtained artificially.
Does not react with nitric acid glass, fluoroplastic-4.
Historical information
The method of obtaining dilute nitric acid by dry distillation of saltpeter with alum and copper sulfate was apparently first described in the treatises of Jabir (Geber in Latinized translations) in the 8th century. This method with certain modifications, the most significant of which was the replacement copper sulfate iron, was used in European and Arab alchemy until the 17th century.
IN XVII century Glauber proposed a method for producing volatile acids by reacting their salts with concentrated sulfuric acid, including nitric acid from potassium nitrate, which made it possible to introduce concentrated nitric acid into chemical practice and study its properties. Method Glauber was used before XX century, and its only significant modification was the replacement of potassium nitrate with cheaper sodium (Chilean) nitrate.
During the time of M.V. Lomonosov, nitric acid was called strong vodka. Industrial production, application and effect on the body
Nitric acid production
Nitric acid is one of the largest volume products chemical industry.
Nitric acid production
The modern method of its production is based on the catalytic oxidation of synthetic ammonia on platinum-rhodium catalysts(process Ostwald) to the mixture oxides nitrogen(nitrous gases), with their further absorption water
Concentration The amount of nitric acid obtained by this method varies depending on the technological design of the process from 45 to 58%. Alchemists were the first to obtain nitric acid by heating a mixture of saltpeter and iron sulfate:
4KNO 3 + 2(FeSO 4 · 7H 2 O)(t°) > Fe2O3 + 2K2SO4+2HNO3^+ NO 2^ + 13H2O
Pure nitric acid was first obtained by Johann Rudolf Glauber by treating nitrate with concentrated sulfuric acid:
KNO 3 + H2SO4(conc.) (t°) > KHSO 4+HNO3^
By further distillation the so-called "smoking Nitric acid", containing virtually no water.
With oxidation states +1, +2, +3, +4, +5.
The oxides N20 and N0 are non-salt-forming (what does this mean?), and the remaining oxides are acidic: N2O3 corresponds to nitrous acid HN02, and N205 corresponds to nitric acid HNO3. Nitrogen oxide (IV) NO2, when dissolved in water, simultaneously forms two acids - HNO2 and HNO3.
If it dissolves in water in the presence of excess oxygen, only nitric acid is obtained
4N02 + 02 + 2H20 = 4HNO3
Nitrogen oxide (IV) NO2 is a brown, very poisonous gas. It is easily obtained by the oxidation of colorless, non-salt-forming nitric oxide (N) by atmospheric oxygen:
Lesson content lesson notes supporting frame lesson presentation acceleration methods interactive technologies Practice tasks and exercises self-test workshops, trainings, cases, quests homework discussion questions rhetorical questions from students Illustrations audio, video clips and multimedia photographs, pictures, graphics, tables, diagrams, humor, anecdotes, jokes, comics, parables, sayings, crosswords, quotes Add-ons abstracts articles tricks for the curious cribs textbooks basic and additional dictionary of terms other Improving textbooks and lessonscorrecting errors in the textbook updating a fragment in a textbook, elements of innovation in the lesson, replacing outdated knowledge with new ones Only for teachers perfect lessons calendar plan for a year guidelines discussion programs Integrated LessonsIn order to depict the formula of a salt graphically, you should:
1. Write the empirical formula of this compound correctly.
2. Considering that any salt can be represented as a product of neutralization of the corresponding acid and base, the formulas of the acid and base that form this salt should be depicted separately.
For example:
Ca(HSO 4) 2 - calcium hydrogen sulfate can be obtained by incomplete neutralization of sulfuric acid H 2 SO 4 with calcium hydroxide Ca(OH) 2.
3. Determine how many molecules of acid and base are required to obtain a molecule of this salt.
For example:
To obtain a Ca(HSO 4) 2 molecule, one molecule of base (one calcium atom) and two molecules of acid (two acid residues HSO 4 1) are required.
Ca(OH) 2 + 2H 2 SO 4 = Ca(HSO 4) 2 + 2H 2 O.
Next, you should construct graphic images of the formulas of the established number of molecules of the base and acid and, mentally removing the hydroxyl anions of the base and hydrogen cations of the acid that participate in the neutralization reaction and form water, obtain a graphic image of the formula of the salt:
O – H H - O O O O
Ca 
+ → Ca + 2 H - O - H
O – H H - O O O O
H- O O H- O O
Salts are solid crystalline substances. Based on their solubility in water, they can be divided into:
1) highly soluble,
2) slightly soluble,
3) practically insoluble.
Most nitric salts and acetic acid, as well as potassium, sodium and ammonium salts - soluble in water.
Salts have wide range melting temperatures and thermal decomposition.
The chemical properties of salts characterize their relationship to metals, alkalis, acids and salts.
1. Salts in solutions interact with more active metals.
A more active metal replaces a less active metal in the salt (see Appendix Table 9).
For example:
Рb(NO 3) 2 + Zn = Рb + Zn(NO 3) 2,
Hg(NO 3) 2 + Cu = Hg + Cu(NO 3) 2.
2. Salt solutions react with alkalis, in this case a new foundation is obtained and new salt.
For example:
CuSO 4 + 2KOH = Cu(OH) 2 + 2K 2 SO 4,
FeCl 3 + 3NaOH = Fe(OH) 3 + 3NaCl.
3. Salts react with solutions of stronger or less volatile acids, this produces a new salt and a new acid.
For example:
a) as a result of the reaction, a weaker acid or a more volatile acid is formed:
Na 2 S + 2HC1 = 2NaCl + H 2 S
b) reactions of salts of strong acids with weaker acids are also possible if the reaction results in the formation of a slightly soluble salt:
СuSO 4 + Н 2 S = СuS + H 2 SO 4 .
4. Salts in solutions enter into exchange reactions with other salts, this produces two new salts.
For example:
NaС1 + AgNO 3 = AgCl + NaNO 3,
CaCI 2 + Na 2 CO 3 = CaCO 3 + 2NaCl,
CuSO 4 + Na 2 S = CuS+ Na 2 SO 4.
It should be remembered that exchange reactions proceed almost to completion if one of the reaction products is released from the reaction sphere in the form of a precipitate, gas, or if water or other weak electrolyte is formed during the reaction.
Oxides. Nitrogen forms five oxides with oxidation states +1, +2, +3, +4, +5.
The oxides N 2 O and NO are non-salt-forming (what does this mean?), and the remaining oxides are acidic: corresponds to nitrous acid, a - nitric acid. Nitrogen oxide (IV), when dissolved in water, simultaneously forms two acids - HNO 2 and HNO 3:
2NO 2 + H 2 O = HNO 2 + HNO 3.
If it is dissolved in water in the presence of excess oxygen, only nitric acid is obtained:
4NO 2 + O 2 + 2H 2 O = 4HNO 3.
Nitrogen oxide (IV) NO 2 is a brown, very poisonous gas. It is easily obtained by the oxidation of colorless, non-salt-forming nitrogen oxide (II) with air oxygen:
2NO + O 2 = 2NO 2.
Nitric acid HNO 3. It is a colorless liquid that “smoke” in air. When stored in the light, concentrated nitric acid turns yellow, as it partially decomposes to form brown gas NO 2:
4HNO 3 = 2H 2 O + 4NO 2 + O 2.
Nitric acid exhibits all the typical properties of strong acids: it interacts with metal oxides and hydroxides, with salts (make up the appropriate reaction equations).
Laboratory experiment No. 32
Properties of dilute nitric acid
Carry out experiments to prove that nitric acid exhibits the typical properties of acids.
|
Nitric acid behaves in a special way with metals - none of the metals displaces hydrogen from nitric acid, regardless of its concentration (for sulfuric acid this behavior is characteristic only in its concentrated state). This is explained by the fact that HNO 3 is a strong oxidizing agent; in it, nitrogen has a maximum oxidation state of +5. It is this that will be restored when interacting with metals.
The reduction product depends on the position of the metal in the stress series, on the acid concentration and on the reaction conditions. For example, when reacting with copper, concentrated nitric acid is reduced to nitric oxide (IV):
Laboratory experiment No. 33
Reaction of concentrated nitric acid with copper
| Carefully pour 1 ml of concentrated nitric acid into the test tube. Using the tip of a glass tube, scoop up a little copper powder and pour it into a test tube with acid. (If there is no copper powder in your office, you can use a small piece of very thin copper wire, which must first be rolled into a ball.) What do you observe? Why does the reaction occur without heating? Why does this experiment not require the use of a fume hood? If the area of contact between copper and nitric acid is less than the proposed experimental option, then what conditions must be observed? After the experiment, immediately place the test tubes with their contents in a fume hood. Write down the reaction equation and consider redox processes. |
Iron and aluminum, when exposed to concentrated HNO 2, are covered with a durable oxide film, which protects the metal from further oxidation, i.e. the acid passivates the metals. Therefore, nitric acid, like sulfuric acid, can be transported in steel and aluminum tanks.
Nitric acid oxidizes many organic substances and discolors dyes. This usually releases a lot of heat and the substance ignites. So, if a drop of turpentine is added to nitric acid, a bright flash occurs, and a smoldering splinter in the nitric acid lights up (Fig. 135).
Rice. 135.
Burning a splinter in nitric acid
Nitric acid is widely used in the chemical industry for the production nitrogen fertilizers, plastics, artificial fibers, organic dyes and varnishes, medicinal and explosives (Fig. 136).
Rice. 136.
Nitric acid is used to produce:
1 - fertilizers; 2 - plastics; 3 - medicines; 4 - varnishes; 5 - artificial fibers; 6 - explosives
Nitric acid salts - nitrates are obtained by the action of acid on metals, their oxides and hydroxides. Sodium, potassium, calcium and ammonium nitrates are called nitrates: NaNO 3 - sodium nitrate, KNO 3 - potassium nitrate, Ca(NO 3) 2 - calcium nitrate, NH 4 NO 3 - ammonium nitrate. Nitrate is used as nitrogen fertilizer.
Potassium nitrate is also used in the manufacture of black gunpowder, and from ammonium nitrate, as you already know, they are preparing the explosive ammonal. Silver nitrate, or lapis, AgNO 3 is used in medicine as a cauterizing agent.
Almost all nitrates are highly soluble in water. When heated, they decompose releasing oxygen, for example:
Nitrous acid exists either in solution or in the gas phase. It is unstable and disintegrates into vapors when heated:
2HNO 2 “NO+NO 2 +H 2 O
Aqueous solutions of this acid decompose when heated:
3HNO 2 “HNO 3 +H 2 O+2NO
This reaction is reversible, therefore, although the dissolution of NO 2 is accompanied by the formation of two acids: 2NO 2 + H 2 O = HNO 2 + HNO 3
Practically, by reacting NO 2 with water, HNO 3 is obtained:
3NO 2 +H 2 O=2HNO 3 +NO
In terms of acidic properties, nitrous acid is only slightly stronger than acetic acid. Its salts are called nitrites and, unlike the acid itself, are stable. From solutions of its salts, a solution of HNO 2 can be obtained by adding sulfuric acid:
Ba(NO 2) 2 +H 2 SO 4 =2HNO 2 +BaSO 4 ¯
Based on data on its compounds, two types of structure of nitrous acid are suggested:
which correspond to nitrites and nitro compounds. Nitrites active metals have a type I structure, and low-active metals have a type II structure. Almost all salts of this acid are highly soluble, but silver nitrite is the most difficult. All salts of nitrous acid are poisonous. For chemical technology, KNO 2 and NaNO 2 are important, which are necessary for the production of organic dyes. Both salts are obtained from nitrogen oxides:
NO+NO 2 +NaOH=2NaNO 2 +H 2 O or when heating their nitrates:
KNO 3 +Pb=KNO 2 +PbO
Pb is necessary to bind the released oxygen.
Of the chemical properties of HNO 2, oxidative properties are more pronounced, while it itself is reduced to NO:
However, many examples of such reactions can be given where nitrous acid exhibits restorative properties:
The presence of nitrous acid and its salts in a solution can be determined by adding a solution of potassium iodide and starch. Nitrite ion oxidizes iodine anion. This reaction requires the presence of H +, i.e. occurs in an acidic environment.
Nitric acid
In laboratory conditions, nitric acid can be obtained by the action of concentrated sulfuric acid on nitrates:
NaNO 3 +H 2 SO 4(k) =NaHSO 4 +HNO 3 The reaction occurs with low heating.
The production of nitric acid on an industrial scale is carried out by the catalytic oxidation of ammonia with atmospheric oxygen:
1. First, a mixture of ammonia and air is passed over a platinum catalyst at 800°C. Ammonia is oxidized to nitric oxide (II):
4NH 3 + 5O 2 =4NO+6H 2 O
2. Upon cooling, further oxidation of NO occurs to NO 2: 2NO+O 2 =2NO 2
3. The resulting nitrogen oxide (IV) dissolves in water in the presence of excess O 2 to form HNO 3: 4NO 2 +2H 2 O+O 2 =4HNO 3
The starting products - ammonia and air - are thoroughly cleaned of harmful impurities that poison the catalyst (hydrogen sulfide, dust, oils, etc.).
The resulting acid is dilute (40-60% acid). Concentrated nitric acid (96-98% strength) is obtained by distilling dilute acid in a mixture with concentrated sulfuric acid. In this case, only nitric acid evaporates.
Nitric acid is a colorless liquid with a pungent odor. Very hygroscopic, “smoke” in air, because its vapors with air moisture form drops of fog. Mixes with water in any ratio. At -41.6°C it goes into a crystalline state. Boils at 82.6°C.
In HNO 3, the valency of nitrogen is 4, the oxidation state is +5. The structural formula of nitric acid is depicted as follows:
Both oxygen atoms, bound only to nitrogen, are equivalent: they are at the same distance from the nitrogen atom and each carry half the charge of an electron, i.e. the fourth part of nitrogen is divided equally between two oxygen atoms.
The electronic structure of nitric acid can be deduced as follows:
1. A hydrogen atom bonds with an oxygen atom by a covalent bond:
2. Due to the unpaired electron, the oxygen atom forms a covalent bond with the nitrogen atom:
3. Two unpaired electrons of the nitrogen atom form a covalent bond with the second oxygen atom:
4. The third oxygen atom, when excited, forms a free 2p- orbital by electron pairing. The interaction of a nitrogen lone pair with a vacant orbital of the third oxygen atom leads to the formation of a nitric acid molecule:
Chemical properties
1. Dilute nitric acid exhibits all the properties of acids. It belongs to strong acids. Dissociates in aqueous solutions:
HNO 3 “Н + +NO - 3 Partially decomposes under the influence of heat and light:
4HNO 3 =4NO 2 +2H 2 O+O 2 Therefore, store it in a cool and dark place.
2. Nitric acid is characterized exclusively by oxidizing properties. The most important chemical property is interaction with almost all metals. Hydrogen is never released. The reduction of nitric acid depends on its concentration and the nature of the reducing agent. The degree of oxidation of nitrogen in the reduction products ranges from +4 to -3:
HN +5 O 3 ®N +4 O 2 ®HN +3 O 2 ®N +2 O®N +1 2 O®N 0 2 ®N -3 H 4 NO 3
The reduction products from the interaction of nitric acid of different concentrations with metals of different activity are shown in the diagram below.
Concentrated nitric acid at ordinary temperatures does not interact with aluminum, chromium, and iron. It puts them into a passive state. A film of oxides forms on the surface, which is impermeable to concentrated acid.
3. Nitric acid does not react with Pt, Rh, Ir, Ta, Au. Platinum and gold dissolve in “royal vodka” - a mixture of 3 volumes of concentrated of hydrochloric acid and 1 volume of concentrated nitric acid:
Au+HNO 3 +3HCl= AuCl 3 +NO+2H 2 O HCl+AuCl 3 =H
3Pt+4HNO 3 +12HCl=3PtCl 4 +4NO+8H 2 O 2HCl+PtCl 4 =H 2
The effect of “regia vodka” is that nitric acid oxidizes hydrochloric acid to free chlorine:
HNO 3 +HCl=Cl 2 +2H 2 O+NOCl 2NOCl=2NO+Cl 2 The released chlorine combines with metals.
4. Non-metals are oxidized with nitric acid to the corresponding acids, and depending on the concentration it is reduced to NO or NO 2:
S+bHNO 3(conc) =H 2 SO 4 +6NO 2 +2H 2 OP+5HNO 3(conc) =H 3 PO 4 +5NO 2 +H 2 O I 2 +10HNO 3(conc) =2HIO 3 +10NO 2 +4H 2 O 3P+5HNO 3(p asb) +2H 2 O= 3H 3 PO 4 +5NO
5. It also interacts with organic compounds.
Salts of nitric acid are called nitrates and are crystalline substances that are highly soluble in water. They are obtained by the action of HNO 3 on metals, their oxides and hydroxides. Potassium, sodium, ammonium and calcium nitrates are called nitrates. Nitrate is used mainly as mineral nitrogen fertilizers. In addition, KNO 3 is used to prepare black powder (a mixture of 75% KNO 3, 15% C and 10% S). The explosive ammonal is made from NH 4 NO 3, aluminum powder and trinitrotoluene.
Salts of nitric acid decompose when heated, and the decomposition products depend on the position of the salt-forming metal in the series of standard electrode potentials:
Decomposition when heated (thermolysis) - important property salts of nitric acid.
2KNO 3 =2KNO 2 +O 2
2Cu(NO 3) 2 = 2CuO+NO 2 +O 2
Salts of metals located in the series to the left of Mg form nitrites and oxygen, from Mg to Cu - metal oxide, NO 2 and oxygen, after Cu - free metal, NO 2 and oxygen.
Application
Nitric acid is the most important product of the chemical industry. Large quantities are spent on the preparation of nitrogen fertilizers, explosives, dyes, plastics, artificial fibers and other materials. Smoking
Nitric acid is used in rocket technology as a rocket fuel oxidizer.
