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    • The maximum and minimum oxidation state of nitrogen. Nitrogen and its compounds

    20.09.2019

    To place correctly oxidation states There are four rules to keep in mind.

    1) In a simple substance, the oxidation state of any element is 0. Examples: Na 0, H 0 2, P 0 4.

    2) You should remember the elements for which are characteristic constant oxidation states. All of them are listed in the table.


    3) The highest oxidation state of an element, as a rule, coincides with the number of the group in which this element is located (for example, phosphorus is in group V, the highest SD of phosphorus is +5). Important exceptions: F, O.

    4) The search for the oxidation states of the remaining elements is based on a simple rule:

    In a neutral molecule, the sum of the oxidation states of all elements is equal to zero, and in an ion - the charge of the ion.

    A few simple examples for determining oxidation states

    Example 1. It is necessary to find the oxidation states of elements in ammonia (NH 3).

    Solution. We already know (see 2) that Art. OK. hydrogen is +1. It remains to find this characteristic for nitrogen. Let x be the desired oxidation state. We compose the simplest equation: x + 3 (+1) \u003d 0. The solution is obvious: x \u003d -3. Answer: N -3 H 3 +1.


    Example 2. Specify the oxidation states of all atoms in the H 2 SO 4 molecule.

    Solution. The oxidation states of hydrogen and oxygen are already known: H(+1) and O(-2). We compose an equation for determining the degree of oxidation of sulfur: 2 (+1) + x + 4 (-2) \u003d 0. Solving this equation, we find: x \u003d +6. Answer: H +1 2 S +6 O -2 4 .


    Example 3. Calculate the oxidation states of all elements in the Al(NO 3) 3 molecule.

    Solution. The algorithm remains unchanged. The composition of the "molecule" of aluminum nitrate includes one atom of Al (+3), 9 oxygen atoms (-2) and 3 nitrogen atoms, the oxidation state of which we have to calculate. Corresponding equation: 1 (+3) + 3x + 9 (-2) = 0. Answer: Al +3 (N +5 O -2 3) 3.


    Example 4. Determine the oxidation states of all atoms in the (AsO 4) 3- ion.

    Solution. In this case, the sum of the oxidation states will no longer be equal to zero, but to the charge of the ion, i.e., -3. Equation: x + 4 (-2) = -3. Answer: As(+5), O(-2).

    What to do if the oxidation states of two elements are unknown

    Is it possible to determine the oxidation states of several elements at once using a similar equation? If we consider this problem from the point of view of mathematics, the answer will be negative. A linear equation with two variables cannot have a unique solution. But we are not just solving an equation!

    Example 5. Determine the oxidation states of all elements in (NH 4) 2 SO 4.

    Solution. The oxidation states of hydrogen and oxygen are known, but sulfur and nitrogen are not. A classic example of a problem with two unknowns! We will consider ammonium sulfate not as a single "molecule", but as a combination of two ions: NH 4 + and SO 4 2-. We know the charges of ions, each of them contains only one atom with an unknown degree of oxidation. Using the experience gained in solving previous problems, we can easily find the oxidation states of nitrogen and sulfur. Answer: (N -3 H 4 +1) 2 S +6 O 4 -2.

    Conclusion: if the molecule contains several atoms with unknown oxidation states, try to "split" the molecule into several parts.

    How to arrange oxidation states in organic compounds

    Example 6. Indicate the oxidation states of all elements in CH 3 CH 2 OH.

    Solution. Finding oxidation states in organic compounds has its own specifics. In particular, it is necessary to separately find the oxidation states for each carbon atom. You can reason as follows. Consider, for example, the carbon atom in the methyl group. This C atom is connected to 3 hydrogen atoms and an adjacent carbon atom. On the C-H bond, the electron density shifts towards the carbon atom (because the electronegativity of C exceeds the EO of hydrogen). If this displacement were complete, the carbon atom would acquire a charge of -3.

    The C atom in the -CH 2 OH group is bonded to two hydrogen atoms (electron density shift towards C), one oxygen atom (electron density shift towards O) and one carbon atom (we can assume that the shifts in electron density in this case not happening). The oxidation state of carbon is -2 +1 +0 = -1.

    Answer: C -3 H +1 3 C -1 H +1 2 O -2 H +1.

    Do not confuse the concepts of "valence" and "oxidation state"!

    Oxidation state is often confused with valence. Don't make that mistake. I will list the main differences:

    • the oxidation state has a sign (+ or -), valence - no;
    • the degree of oxidation can be equal to zero even in a complex substance, the equality of valency to zero means, as a rule, that the atom of this element is not connected to other atoms (we will not discuss any kind of inclusion compounds and other "exotics" here);
    • the degree of oxidation is a formal concept that acquires real meaning only in compounds with ionic bonds, the concept of "valence", on the contrary, is most conveniently applied in relation to covalent compounds.

    The oxidation state (more precisely, its modulus) is often numerically equal to the valence, but even more often these values ​​do NOT coincide. For example, the oxidation state of carbon in CO 2 is +4; valency C is also equal to IV. But in methanol (CH 3 OH), the valency of carbon remains the same, and the oxidation state of C is -1.

    A small test on the topic "The degree of oxidation"

    Take a few minutes to check how you have understood this topic. You need to answer five simple questions. Good luck!

    VA subgroup is formed by p-elements: nitrogenN, phosphorus

    P, arsenicAs, antimony Sb and bismuth Bi.

    Elements N, P are typical non-metals,

    for nonmetals As and Sb some properties appear

    inherent in metals, bismuth has metallic properties

    predominate, although it is not a typical metal.

    The general formula for valence electrons in elemental

    of the VA group –ns 2 np 3.

    throne. With three unpaired electrons all elements in simple substances form three covalent bonds, but in nitrogen, three bonds unite 2 atoms, forming a very strong

    molecule N N, and for other elements, each atom is bonded to three others to form molecules of the E4 type (be-

    blue phosphorus and yellow arsenic) or polymeric structures.

    In nitrogen, a simple substance in any state of aggregation consists of individual molecules , under normal conditions it is a gas. All other elements are simple substances

    - solid.

    The oxidation state (–3) for elements of the VA group is minimal. It is most stable at N, at

    transition to Bi with an increase in the number of electron layers, its stability pa-

    gives. The elements N, P, As, Sb with hydrogen form EN3 type hydrides,

    exhibiting basic properties, they are most pronounced in ammonia

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    like NH3. In the subgroup, the stability of EN3 compounds and their main properties

    wa decrease.

    All elements of the VA-group exhibit the highest oxidation state +5.

    All of them form oxides of the E2 O5 type (Bi 2 O 5 oxide is unstable), which correspond to acids, the strength of acids weakens when moving down the sub-

    +5 oxidation state is most stable in P . Bi(+5) compounds –

    very strong oxidizing agents. Nitric acid, especially concentrated, exhibits strong oxidizing properties.

    Bismuth has a more stable oxidation state (+3), which is also quite stable for Sb and As. N(+3) compounds, and especially

    P(+3), exhibit strong reducing properties.

    In the oxidation state +3, all elements of the VA group form oxides

    type E 2 O 3. Oxides N and P correspond to weak acids. Oxides and hydroxy

    dys As and Sb are amphoteric, the main character is dominated by the oxide and hydroxy-

    yes Bi(+3). In this way, in the subgroup the acidic nature of oxides and hyd-

    roxides of elements in the oxidation state (+3) weakens, and increase

    basic properties more characteristic of metal hydroxides.

    Elements of the VA group, in addition to the listed oxidation states

    5, +3, –3 also exhibit other intermediate oxidation states.

    For nitrogen, all oxidation states from –1 to +5 are known.

    Nitrogen, like all elements of the second period, differs significantly from its electronic counterparts. . For this reason, as well as due to a large number of oxidation states and a variety of compounds, nitrogen chemistry is considered

    vaetsya separately from other elements of the VA-subgroup.

    The most common element of the VA group in nature is

    there is phosphorus. Its content in the earth's crust is 0.09 wt. %; phosphorus is found

    mainly in the form of calcium phosphate. Nitrogen content - 0.03%, os-

    its new share is concentrated in the atmosphere in the form of N2. Nitrogen content in

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    air by volume is ~ 78%. In very small quantities in the earth

    nitrates of sodium and potassium (nitrate) are found in the bark. Arsenic, antimony and bismuth are rare elements with a content in the earth's crust of 10–5 5. 10–

    four %; in nature, they are found mainly in the form of sulfides.

    Nitrogen and phosphorus are very important elements of the biosphere, therefore, significant

    the main part of nitrates and phosphates produced in the chemical industry

    fats are used as fertilizers, which are necessary for life

    plant vigor. In the human body, N and P play an important role, - nitrogen

    is part of the amino acids that are an integral part of proteins, phosphorus in

    form Ca5 [(PO4 )3 OH] is part of the bones. In the human body are

    averages about 1.8 kg N.

    Some characteristics of the atoms of the elements of the VA group are given in

    The most important characteristics of the atoms of the elements of the VA group

    Electric

    denigrator-

    ness (according to

    atom, nm

    Polling)

    an increase in the number of

    throne layers;

    an increase in the size of an atom;

    decrease in ion energy

    decrease in electrical

    values;

    For comparison, the electronegativity H is 2.2; O - 3.44.

    Nitrogen differs from other elements of the subgroup in a very small orbit

    tal radius and high electronegativity, N is the third in electoral

    trinegativity element, after F and O.

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    Valence electrons N –2s2 2p3 .

    N 2s

    Nitrogen, like other elements of the second period,

    differs markedly from the elements of its subgroup:

    the N atom has only 4 valence orbitals and in compounds it can form

    call only 4 covalent bonds;

    due to the very small atomic radius, nitrogen forms very strong

    a simple substance in any state of aggregation consists of separate

    very strong molecules N

    N and is highly inert;

    in electronegativity, N is second only to F and O;

    nitrogen exhibits all possible oxidation states: -3, -2, -1, 0, +1, +2, +3, +4, +5.

    A large number of oxidation states and a variety of compounds makes

    The chemistry of nitrogen is highly complex. The complexity is also exacerbated by the kinetic difficulties characteristic of many redox reactions.

    due to very strong multiple bonds between atoms

    N and N and O atoms. Therefore, electrode potentials are of little help in determining

    division of OVR products.

    The most stable compound N is a simple substance.

    In aqueous solutions, especially acidic ones, the NH4 + ion is very stable.

    Nitrogen is an integral part of the air, from which N 2 is obtained.

    The main amount of N2 is used for the synthesis of ammonia, from which other nitrogen compounds are then obtained. Among nitrogen compounds, ammonia, nitric acid and their salts find the widest practical application..

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    The annual world production of NH3 is ~ 97 million tons / year, nitric acid

    slots - 27 million tons / year. The chemistry of these most important N compounds will be considered

    ren in the first place, after discussing the properties of a simple substance.

    simple substance

    The N2 molecule is the strongest of all diatomic molecules of simple substances. Three common electron pairs in the N N molecule are located on the bond

    inviting orbitals, there are no electrons in loosening orbitals - this is

    leads to a very high chemical bond energy - 944 kJ / mol (for comparison,

    the binding energy in the O2 molecule is – 495 kJ/mol). Strong bond causes high inertness of molecular nitrogen. The name of this element is associated with the chemical inertness of nitrogen. In Greek, "nitrogen" means

    tea "lifeless".

    Under normal conditions, N2 is a colorless, odorless and tasteless gas.

    The boiling and melting points of N2 are close: –196О С, and –210О С.

    Nitrogen is obtained by fractional distillation of air , - for this air

    at low temperatures, they liquefy, and then begin to raise the temperature.

    Of the components of air, nitrogen has the lowest boiling point and

    forms the lightest fraction. In fractional distillation, one

    temporarily receive oxygen and inert gases.

    The main amount of N2 goes to the production of ammonia, in addition,

    nitrogen is used to create an inert atmosphere, including in the production of

    some metals; liquid nitrogen is also used as a cooling

    giving agent in the laboratory and in industry.

    At room temperature, nitrogen slowly reacts only with Li to form

    formation of Li3 N. When burning in air, magnesium, together with MgO oxide, forms

    Mg3 N2 is also present.

    Nitrides. Binary compounds of nitrogen with elements less electrified

    three-negative than N are called nitrides.

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    Ionic nitrides contain anion N3–. Ionic nitrides form Li,

    metals II and IB-groups; in aqueous solutions, they undergo irreversible

    mu hydrolysis.

    Mg3 N2 + 6H2 O = 2NH3 + 3 Mg(OH)2

    With p-block metals and some light non-metals, nitrogen is ob-

    forms covalent nitrides, for example, AlN, BN.

    Most d-metals form non-stoichiometric interstitial products with nitrogen at high temperatures, in which N atoms occupy spaces.

    thots in crystal lattices of metals. Therefore, such nitrides externally

    its appearance, in terms of electrical and thermal conductivity resemble metals, but differ

    They are distinguished by high chemical inertness, hardness and refractoriness.

    For example, non-stoichiometric Ta and Ti nitrides melt at temperatures above 3200°C.

    Nitrogen does not directly react with halogens, but interacts with oxygen only under extreme conditions(with electric

    discharge).

    The most important in practical terms is the reaction of nitrogen with H2, which results in ammonia.

    N 2 + 3H 2  2NH 3; H0 = –92 kJ/mol.

    The exothermicity of this reaction indicates that the total bond strength in ammonia molecules is higher than in the initial molecules. An increase in temperature, in accordance with the Le Chatelier principle, leads to a shift in the equilibrium towards an endothermic reaction, i.e. in the direction of ammonia decomposition. However, under normal conditions, the reaction is extremely slow.

    however, the activation energy required to weaken strong bonds in nitrogen and hydrogen molecules is too high. Therefore, the process has to be carried out at a temperature of about 5000 C. To shift the equilibrium at high temperature to the right, increase the pressure to 300 - 500 atm., while equilibrium

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    this shifts in the direction of the reaction proceeding with a decrease in the number of gas molecules, i.e. towards the formation of ammonia. The increase in speed is achieved through the use of catalysts. Efficient fused catalyst on os-

    new Fe3 O4 with additions of Al2 O3 and SiO2 and a catalyst based on a metal

    Fe. The synthesis of ammonia from nitrogen and hydrogen is the most important reaction

    industrial chemistry of nitrogen.

    Nitrogen compounds

    Ammonia and ammonium salts

    Nitrogen in ammonia and ammonium salts is in the minimum oxidation state (–3). The oxidation state (–3) is fairly stable with nitrogen.

    Ammonia under normal conditions is a colorless gas with a characteristic

    pungent odor, familiar by the smell of "ammonia" (10% so-

    ammonia in water). This gas is lighter than air, so it can be collected in upside down vessels. Ammonia easily passes into a liquid. To do this, it is enough to cool it at ordinary pressure to -33.5 ° C. The same effect

    effect can be achieved at room temperature, but by increasing the pressure to

    7 - 8 atm. At elevated pressure, liquid ammonia is stored in steel balls

    nah. Evaporating, liquid ammonia causes cooling in the environment. This is the basis for its use in refrigeration. The easy liquefaction of ammonia is due to hydrogen bonds between its molecules. The strength of hydrogen bonds between ammonia molecules is due to the very high electronegativity of nitrogen.

    Liquid ammonia is colorless, undergoes autoprotolysis:

    2NH3  NH4 + + NH2 –

    The constant of this equilibrium is equal to 2 . 10–23 (at –50°C). Liquid ammonia

    is a good ionizing solvent . Ammonium salts and weak

    acids, for example, H2 S, dissolved in liquid ammonia, become strong

    mi acids.

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    Ammonia is highly soluble in water. The high solubility of ammonia in water (up to 700 volumes of NH3 in one volume of water) is also explained by the formation

    We eat hydrogen bonds, but already with water molecules. Concentrated dis-

    the solution contains 25 mass% ammonia and has a density of 0.91 g/cm3. The molar concentration of NH3 in concentrated aqueous solutions reaches ~13

    The NH3 molecule has a pyramidal structure, which is explained by sp3 -

    hybridization of valence atomic orbitals of nitrogen. One of the vertices of the tetrahedron

    ra is occupied by a lone pair of electrons. The N–H bond is quite strong,

    the bond energy is 389 kJ/mol, the bond length is 0.1 nm, the angle between the bonds

    zyami -108.3o. When the H+ cation is attached due to the unshared electron

    pair N, a very stable tetrahedral ammonium ion is formed

    NH4+.

    The presence of an unshared electron pair at N in the NH3 molecule,

    glorifies many characteristic properties of ammonia.

    The NH3 molecule is a good electron pair donor (DEP),

    those. Lewis base, and a very good proton acceptor A(H+),

    those. basis according to Bronsted:

    NH3 + H+  NH4 + . NH3 accepts a proton, like OH– ions: OH– + H+  H2 O

    The acceptor properties of NH3 are weaker than those of the OH– anion. The protolysis constant for NH3 is 1.8. 109 , and for the OH– ion – 1014 .

    Reactions with acids are the most typical reactions for NH3.

    The ability of ammonia to form donor-acceptor bonds on-

    so large that it can tear off hydrogen ions from such a strong bond

    unity like water.

    NH3 + H–– OH  NH4 + ), and the amount of NH4 + and OH– products is small compared to the equilibrium concentration of ammonia. Aqueous solutions of ammonia behave like weak bases. According to the established tradition, ammonia is often designated

    tea with the formula NH4 OH and is called ammonium hydroxide, however, the molecules

    NH4 OH is not present in the solution. The alkaline reaction of an aqueous solution of NH3 is often described

    are not described by the above equilibrium, but as a dissociation of molecules

    NH4OH:

    NH4 OH NH4 + + OH–

    The constant of this equilibrium is 1.8. 10–5 . In one liter, one molar

    th ammonia solution, the concentration of NH4 + and OH– ions is 3.9. 10–3

    mol/l, pH = 11.6.

    The equilibrium between ammonia and OH– can strongly shift to the right the cations of some metals, which form insoluble hydroxides with OH– ions.

    FeCl3 + 3NH3 + 3H–OH  Fe(OH)3  + 3NH4 Cl.

    Ammonia can be used to make insoluble bases.

    Under the action of acids on aqueous solutions of ammonia, ammonium salts are formed.

    NH3 + HCl = NH4Cl

    Almost all ammonium salts are colorless and soluble in water.

    The equilibrium NH3 + H+  NH4 + is strongly shifted to the right (K = 1.8.109),

    this means that, NH3 is a strong proton acceptor, and the NH 4 + cation

    is a weak donor of H+ , i.e. acid according to Bronsted. When alkali is added to ammonium salts, ammonia is formed, which is easy to determine by

    NH4Cl + NaOH = NH3 + H2O + NaCl.

    This reaction is commonly used to detect ammonium ions in solution.

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    Similar reactions can be used for laboratory preparation

    NH3.

    Ammonium chloride (called ammonia) at high temperatures reacts with oxides on the surface of metals, like an acid, exposing pure metal. The use of the solid salt NH4 Cl in soldering metals is also based on this. The "acidic" H+ from the NH4+ ion is able to oxidize very active metals such as Mg.

    Mg + 2NH4 Cl = H2 + MgCl2 + 2NH3

    A characteristic property of ammonium salts is their thermal instability.

    durability. When heated, they decompose quite easily. Products

    The positions are determined by the properties of the acid anion. If the anion exhibits oxidizing properties, then NH4 + is oxidized and the oxidizing anion is reduced.

    NH4 NO2 = N2 + 2H2O

    NH4 NO3 = N2 O + 2H2 O or 2NH4 NO3 = N2 + O2 + 4H2 O

    (NH4)2 Cr2 O7 = N2 + Cr2 O3 + 4H2 O

    From the salts of volatile acids, ammonia and acid (or its anhydrous) are released.

    reed), and in the case of non-volatile acids (for example, H3 PO4) - only NH3. NH4 HCO3 = NH3 + H2 O + CO2

    Ammonium bicarbonate NH4 HCO3 is used in baking

    The resulting gases give the dough the necessary porosity.

    Ammonium salts are used in the manufacture of explosives and in

    as nitrogen fertilizer. Ammonal, used in the practice of blasting, is a mixture of NH4 NO3 salt (72%), Al powder (25%) and coal

    la (3%). This mixture explodes only after detonation.

    The second type of reactions in which NH3 exhibits the properties of an electron donor

    throne pair is formation of amine complexes. Ammonia, as a ligand, joins the cations of many d-elements, forming a chemical

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    NITROGEN N (nitrogenium) , chemical element (at. number 7) VA subgroup of the Periodic Table of Elements. The Earth's atmosphere contains 78% (vol.) nitrogen. To show how large these reserves of nitrogen are, we note that in the atmosphere above each square kilometer of the earth's surface there is so much nitrogen that up to 50 million tons of sodium nitrate or 10 million tons of ammonia (a combination of nitrogen with hydrogen) can be obtained from it, and yet this is a small fraction of the nitrogen contained in the earth's crust. The existence of free nitrogen indicates its inertness and the difficulty of interacting with other elements at ordinary temperatures. Bound nitrogen is part of both organic and inorganic matter. Plant and animal life contains nitrogen bound to carbon and oxygen in proteins. In addition, nitrogen-containing inorganic compounds such as nitrates (NO 3 ), nitrites (NO 2 ), cyanides (CN ), nitrides (N 3 ) and azides (N 3 ). History reference. The experiments of A. Lavoisier, devoted to the study of the role of the atmosphere in maintaining life and combustion processes, confirmed the existence of a relatively inert substance in the atmosphere. Not having established the elemental nature of the gas remaining after combustion, Lavoisier called it azote, which in ancient Greek means "lifeless". In 1772, D. Rutherford from Edinburgh established that this gas is an element and called it "harmful air". The Latin name for nitrogen comes from the Greek words nitron and gen, which means "forming saltpetre".Nitrogen fixation and the nitrogen cycle. The term "nitrogen fixation" refers to the process of fixing atmospheric nitrogen N 2 . In nature, this can happen in two ways: either legumes, such as peas, clover, and soybeans, accumulate nodules on their roots, in which nitrogen-fixing bacteria convert it into nitrates, or atmospheric nitrogen is oxidized by oxygen under conditions of a lightning discharge. S. Arrhenius found that up to 400 million tons of nitrogen are fixed in this way annually. In the atmosphere, nitrogen oxides combine with rainwater to form nitric and nitrous acids. In addition, it has been established that with rain and snow, approx. 6700 g of nitrogen; reaching the soil, they turn into nitrites and nitrates. Plants use nitrates to form plant proteins. Animals, eating these plants, assimilate the protein substances of plants and turn them into animal proteins. After the death of animals and plants, they decompose, nitrogen compounds turn into ammonia. Ammonia is used in two ways: bacteria that do not form nitrates break it down to elements, releasing nitrogen and hydrogen, and other bacteria form nitrites from it, which are oxidized to nitrates by other bacteria. Thus, the nitrogen cycle in nature, or the nitrogen cycle, occurs.The structure of the nucleus and electron shells. In nature, there are two stable isotopes of nitrogen: with a mass number of 14 (contains 7 protons and 7 neutrons) and withmass number 15 (contains 7 protons and 8 neutrons). Their ratio is 99.635:0.365, so the atomic mass of nitrogen is 14.008. Unstable isotopes of nitrogen 12N, 13N, 16N, 17 N obtained artificially.Schematically the electronic structure of the nitrogen atom is: 1 s 2 2s 2 2p x 1 2py 1 2pz 1 . Therefore, on the outer (second) electron shell there are 5 electrons that can participate in the formation of chemical bonds; nitrogen orbitals can also accept electrons, i.e. it is possible to form compounds with an oxidation state from ( II I) to (V), and they are known. see also ATOM STRUCTURE.Molecular nitrogen. From the definitions of gas density, it was established that the nitrogen molecule is diatomic, i.e. the molecular formula of nitrogen is Nє N (or N 2 ). Two nitrogen atoms have three outer 2p- electrons of each atom form a triple bond:N:::N:, forming electron pairs. Measured interatomic distance N N equals 1.095 Å . As in the case of hydrogen (cm. HYDROGEN), there are nitrogen molecules with different nuclear spins symmetric and antisymmetric. At ordinary temperature, the ratio of symmetric and antisymmetric forms is 2:1. In the solid state, two modifications of nitrogen are known: a cubic and b hexagonal with transition temperature a® b237.39° C. Modification b melts at 209.96° C and boils at 195.78° C at 1 atm ( cm. tab. one). The dissociation energy of a mole (28.016 g or 6.023 Ch 10 23 molecules) of molecular nitrogen into atoms ( N22N) equals approximately 225 kcal. Therefore, atomic nitrogen can be formed in a quiet electrical discharge and is chemically more active than molecular nitrogen.Receipt and application. The method of obtaining elemental nitrogen depends on the required purity. Huge amounts of nitrogen are obtained for the synthesis of ammonia, while small admixtures of noble gases are acceptable.nitrogen from the atmosphere. The economic release of nitrogen from the atmosphere is due to the low cost of the method of liquefying purified air (water vapor, CO 2 , dust, other impurities removed). Successive cycles of compression, cooling and expansion of such air lead to its liquefaction. Liquid air is subjected to fractional distillation with a slow rise in temperature. Noble gases are released first, then nitrogen, and liquid oxygen remains. Purification is achieved by multiple fractionation processes. This method produces many millions of tons of nitrogen annually, mainly for the synthesis of ammonia, which is the feedstock in the technology for the production of various nitrogen-containing compounds for industry and agriculture. In addition, a purified nitrogen atmosphere is often used when the presence of oxygen is unacceptable.laboratory methods. Small amounts of nitrogen can be obtained in the laboratory in various ways by oxidizing ammonia or the ammonium ion, for example:The process of oxidation of the ammonium ion with the nitrite ion is very convenient:Other methods are also known decomposition of azides when heated, decomposition of ammonia with copper (II) oxide, interaction of nitrites with sulfamic acid or urea:With the catalytic decomposition of ammonia at high temperatures, nitrogen can also be obtained: physical properties. Some physical properties of nitrogen are given in Table. one.

    Table 1. SOME PHYSICAL PROPERTIES OF NITROGEN
    Density, g / cm 3 0.808 (liquid)
    Melting point, °С –209,96
    Boiling point, °С –195,8
    Critical temperature, °C –147,1
    Critical pressure, atm a 33,5
    Critical density, g/cm 3 a 0,311
    Specific heat capacity, J / (molCh K) 14.56 (15°C)
    Electronegativity according to Pauling 3
    covalent radius, 0,74
    crystal radius, 1.4 (M 3–)
    Ionization potential, V b
    the first 14,54
    second 29,60
    a The temperature and pressure at which the densityliquid and gaseous nitrogen are the same.
    b The amount of energy required to remove the first outer and following electrons, based on 1 mole of atomic nitrogen.
    Chemical properties. As already noted, the predominant property of nitrogen under normal conditions of temperature and pressure is its inertness, or low chemical activity. The electronic structure of nitrogen contains an electron pair for 2 s-level and three half-filled 2 R-orbitals, so one nitrogen atom can bind no more than four other atoms, i.e. its coordination number is four. The small size of an atom also limits the number of atoms or groups of atoms that can be bonded to it. Therefore, many compounds of other members of the VA subgroup either have no analogues among nitrogen compounds at all, or similar nitrogen compounds turn out to be unstable. Yes, PCl 5 stable connection, and NCl 5 does not exist. A nitrogen atom is able to bond with another nitrogen atom, forming several fairly stable compounds, such as hydrazine N 2 H 4 and metal azides MN 3 . This type of bond is unusual for chemical elements (with the exception of carbon and silicon). At elevated temperatures, nitrogen reacts with many metals to form partially ionic nitrides M x N y . In these compounds, nitrogen is negatively charged. In table. 2 shows the oxidation states and examples of the corresponding compounds.

    Table 2. NITROGEN OXIDATION STATES AND RELATED COMPOUNDS

    Oxidation state

    Connection examples
    Ammonia NH 3, ammonium ion NH 4 +, nitrides M 3 N 2
    Hydrazine N 2 H 4
    Hydroxylamine NH 2 OH
    sodium hyponitrite Na 2 N 2 O 2, nitric oxide (I) N 2 O
    Nitric oxide(II) NO
    Nitric oxide (III) N 2 O 3, sodium nitrite NaNO 2
    Nitric oxide (IV) NO 2, dimer N 2 O 4
    Nitric oxide (V) N 2 O 5 , Nitric acid HNO3 and its salts (nitrates)
    Nitrides. Compounds of nitrogen with more electropositive elements, metals and non-metals nitrides are similar to carbides and hydrides. They can be divided depending on the nature of the MN bond into ionic, covalent and with an intermediate type of bond. As a rule, these are crystalline substances.Ionic nitrides. Bonding in these compounds involves the transfer of electrons from the metal to nitrogen with the formation of the N ion 3 . These nitrides include Li 3 N, Mg 3 N 2, Zn 3 N 2 and Cu 3 N 2 . In addition to lithium, other alkali metals of the IA subgroup do not form nitrides. Ionic nitrides have high melting points, react with water to form NH 3 and metal hydroxides.covalent nitrides. When the electrons of nitrogen participate in the formation of a bond together with the electrons of another element without transferring them from nitrogen to another atom, nitrides with a covalent bond are formed. Hydrogen nitrides (such as ammonia and hydrazine) are fully covalent, as are nitrogen halides (NF 3 and NCl 3 ). Covalent nitrides include, for example, Si 3 N 4 , P 3 N 5 and BN are highly stable white substances, and BN has two allotropic modifications: hexagonal and diamond-like. The latter is formed at high pressures and temperatures and has a hardness close to that of diamond.Nitrides with an intermediate type of bond. Transition elements in reaction with NH 3 at high temperature they form an unusual class of compounds in which nitrogen atoms are distributed between regularly spaced metal atoms. There is no clear displacement of electrons in these compounds. Examples of such nitrides Fe 4 N, W 2 N, Mo 2 N, Mn 3 N 2 . These compounds are generally completely inert and have good electrical conductivity.Hydrogen compounds of nitrogen. Nitrogen and hydrogen interact to form compounds that vaguely resemble hydrocarbons. (see also ORGANIC CHEMISTRY). The stability of hydrogen nitrogens decreases with an increase in the number of nitrogen atoms in the chain, in contrast to hydrocarbons, which are also stable in long chains. The most important hydrogen nitrides ammonia NH 3 and hydrazine N 2 H 4 . These also include hydrazoic acid HNNN (HN 3). Ammonia NH 3 . Ammonia is one of the most important industrial products of the modern economy. At the end of the 20th century The US produced approx. 13 million tons of ammonia annually (in terms of anhydrous ammonia).The structure of the molecule. NH 3 molecule has an almost pyramidal structure. The connection angle HNH is 107° , which is close to the value of the tetrahedral angle 109° . The unshared electron pair is equivalent to the attached group, as a result, the coordination number of nitrogen is 4 and nitrogen is located in the center of the tetrahedron.properties of ammonia. Some physical properties of ammonia in comparison with water are given in table. 3.

    Table 3. SOME PHYSICAL PROPERTIES OF AMMONIA AND WATER

    Property
    Density, g / cm 3 0.65 (-10°C) 1.00 (4.0°C)
    Melting point, °С –77,7 0
    Boiling point, °С –33,35 100
    Critical temperature, °C 132 374
    Critical pressure, atm 112 218
    Enthalpy of vaporization, J/g 1368 (-33°C) 2264 (100°C)
    Enthalpy of melting, J/g 351 (-77°C) 334 (0°C)
    Electrical conductivity 5h 10 -11 (-33°C) 4h 10 -8 (18°C)
    The boiling and melting points of ammonia are much lower than those of water, despite the similarity of molecular weights and the similarity of the structure of the molecules. This is due to the relatively greater strength of intermolecular bonds in water than in ammonia (such an intermolecular bond is called hydrogen).ammonia as a solvent. The high dielectric constant and dipole moment of liquid ammonia make it possible to use it as a solvent for polar or ionic inorganic substances. Ammonia solvent occupies an intermediate position between water and organic solvents such as ethyl alcohol. Alkali and alkaline earth metals dissolve in ammonia, forming dark blue solutions. It can be assumed that solvation and ionization of valence electrons occurs in solution according to the schemeThe blue color is associated with solvation and the movement of electrons or with the mobility of "holes" in a liquid. At a high concentration of sodium in liquid ammonia, the solution takes on a bronze color and is characterized by high electrical conductivity. The unbound alkali metal can be separated from such a solution by evaporation of the ammonia or by adding sodium chloride. Solutions of metals in ammonia are good reducing agents. Autoionization occurs in liquid ammoniasimilar to the process taking place in water:Some chemical properties of both systems are compared in Table. four.

    Liquid ammonia as a solvent has an advantage in some cases where it is impossible to carry out reactions in water due to the rapid interaction of components with water (for example, oxidation and reduction). For example, in liquid ammonia, calcium reacts with KCl to form CaCl 2 and K, since CaCl 2 is insoluble in liquid ammonia, but K is soluble, and the reaction proceeds completely. In water, such a reaction is impossible due to the rapid interaction of Ca with water.

    Getting ammonia. Gaseous NH 3 is released from ammonium salts under the action of a strong base, for example, NaOH:The method is applicable in laboratory conditions. Small ammonia production is also based on the hydrolysis of nitrides, such as Mg 3 N 2 , water. Calcium cyanamide CaCN 2 when interacting with water, it also forms ammonia. The main industrial method for producing ammonia is its catalytic synthesis from atmospheric nitrogen and hydrogen at high temperature and pressure:Hydrogen for this synthesis is obtained by thermal cracking of hydrocarbons, the action of water vapor on coal or iron, the decomposition of alcohols with water vapor, or the electrolysis of water. Many patents have been obtained for the synthesis of ammonia, differing in the process conditions (temperature, pressure, catalyst). There is a method of industrial production during the thermal distillation of coal. The names of F. Haber and K. Bosch are associated with the technological development of ammonia synthesis.

    Table 4. COMPARISON OF REACTIONS IN WATER AND AMMONIA MEDIUM

    Water environment

    Ammonia medium

    Neutralization
    OH - + H 3 O + ® 2H 2 O

    NH 2 - + NH 4 + ® 2NH 3

    Hydrolysis (protolysis)
    PCl 5 + 3H 2 O POCl 3 + 2H 3 O + + 2Cl –

    PCl 5 + 4NH 3 PNCl 2 + 3NH 4 + + 3Cl -

    substitution
    Zn + 2H 3 O + ® Zn 2+ + 2H 2 O + H 2

    Zn + 2NH 4 + ® Zn 2+ + 2NH 3 + H 2

    solvation (complexation )
    Al 2 Cl 6 + 12H 2 O 2 3+ + 6Cl -

    Al 2 Cl 6 + 12NH 3 2 3+ + 6Cl -

    Amphoteric
    Zn 2+ + 2OH - Zn (OH) 2

    Zn 2+ + 2NH 2 - Zn (NH 2) 2
    Zn(OH) 2 + 2H 3 O + Zn 2+ + 4H 2 O

    Zn(NH 2) 2 + 2NH 4 + Zn 2+ + 4NH 3
    Zn(OH) 2 + 2OH – Zn(OH) 4 2–

    Zn(NH 2) 2 + 2NH 2 – Zn(NH 2) 4 2–
    Chemical properties of ammonia. In addition to the reactions mentioned in Table. 4, ammonia reacts with water to form an NH compound 3 H H 2 O, often mistakenly thought to be ammonium hydroxide NH 4 OH; actually existence NH4OH solution has not been proven. An aqueous solution of ammonia ("ammonia") consists mainly of NH 3, H 2 O and small concentrations of NH ions 4+ and OH , formed during dissociationThe main character of ammonia is explained by the presence of a lone electron pair of nitrogen: NH 3 . Therefore, NH 3 this is the Lewis base, which has the highest nucleophilic activity, manifested in the form of association with a proton, or the nucleus of a hydrogen atom:Any ion or molecule capable of accepting an electron pair (electrophilic compound) will interact with NH 3 to form a coordination compound. For example: Symbol M n+ represents a transition metal ion (B-subgroups of the periodic table, for example, Cu 2+ , Mn 2+ andetc.). Any protic (i.e. H-containing) acid reacts with ammonia in aqueous solution to form ammonium salts such as ammonium nitrate NH 4 NO 3 , ammonium chloride NH 4 Cl, ammonium sulfate (NH 4) 2 SO 4 , ammonium phosphate (NH 4) 3PO4 . These salts are widely used in agriculture as fertilizers to introduce nitrogen into the soil. Ammonium nitrate is also used as an inexpensive explosive; for the first time it was applied with fuel oil (diesel oil). An aqueous solution of ammonia is used directly for introduction into the soil or with irrigation water. Urea NH 2 CONH 2 , obtained by synthesis from ammonia and carbon dioxide, is also a fertilizer. Gaseous ammonia reacts with metals such as Na and K to form amides:Ammonia reacts with hydrides and nitrides also to form amides:Alkali metal amides (for example, NaNH 2) react with N 2 O when heated, forming azides: Gaseous NH 3 reduces heavy metal oxides to metals at high temperature, presumably due to the hydrogen formed from the decomposition of ammonia into N 2 and H2: Hydrogen atoms in the NH molecule 3 can be replaced by halogen. Iodine reacts with a concentrated solution of NH 3 , forming a mixture of substances containing N I 3 . This substance is very unstable and explodes at the slightest mechanical impact. In the reaction of NH 3 s Cl 2 chloramines NCl 3 , NHCl 2 and NH 2 Cl are formed. When exposed to ammonia sodium hypochlorite NaOCl (formed from NaOH and Cl2 ) the final product is hydrazine:Hydrazine. The above reactions are a method for preparing hydrazine monohydrate of the composition N 2 H 4 H H 2 O. Anhydrous hydrazine is formed by special distillation of the monohydrate with BaO or other water-removing substances. The properties of hydrazine slightly resemble hydrogen peroxide H 2O2 . Pure Anhydrous Hydrazine– colorless hygroscopic liquid, boiling at 113.5°C ; dissolves well in water, forming a weak base In an acidic environment (H + ) hydrazine forms soluble hydrazonium salts of the type + X . The ease with which hydrazine and some of its derivatives (eg, methylhydrazine) react with oxygen allows it to be used as a component of liquid propellant. Hydrazine and all its derivatives are highly toxic.nitrogen oxides. In compounds with oxygen, nitrogen exhibits all oxidation states, forming oxides: N2 O, NO, N 2 O 3, NO 2 (N 2 O 4), N 2 O 5. Little information is available on the formation of nitrogen peroxides (NO 3 , NO 4). Nitric oxide (I) N 2 O (dianitrogen monoxide) is obtained by thermal dissociation of ammonium nitrate:The molecule has a linear structureO is rather inert at room temperature, but at high temperatures it can support the combustion of readily oxidizing materials. N 2 O, known as "laughing gas", is used for mild anesthesia in medicine.Nitric oxide(II) NO colorless gas, is one of the products of the catalytic thermal dissociation of ammonia in the presence of oxygen:NO is also formed by the thermal decomposition of nitric acid or by the reaction of copper with dilute nitric acid:NO can be obtained by synthesis from simple substances (N 2 and O2 ) at very high temperatures, such as in an electrical discharge. The structure of the NO molecule has one unpaired electron. Compounds with such a structure interact with electric and magnetic fields. In the liquid or solid state, the oxide is blue because the unpaired electron causes partial association in the liquid state and weak dimerization in the solid state: 2NO N 2 O 2 . Nitric oxide (III) N2O3 (nitrogen trioxide) nitrous anhydride: N 2 O 3 + H 2 O 2HNO 2. Pure N 2 O 3 can be obtained as a blue liquid at low temperatures (20° C) from an equimolecular mixture of NO and NO 2. N2O3 stable only in the solid state at low temperatures (m.p. 102.3° C), in the liquid and gaseous state, it decomposes again into NO and NO 2 . Nitric oxide (IV) NO 2 (nitrogen dioxide) also has an unpaired electron in the molecule ( see above nitric(II) oxide). A three-electron bond is assumed in the structure of the molecule, and the molecule exhibits the properties of a free radical (one line corresponds to two paired electrons):obtained by catalytic oxidation of ammonia in excess oxygen or by oxidation of NO in air:as well as reactions:At room temperature NO 2 The gas is dark brown in color, has magnetic properties due to the presence of an unpaired electron. At temperatures below 0°C NO 2 molecule dimerizes to dinitrogen tetroxide, and at 9.3° C dimerization proceeds completely: 2NO2N2O4 . In the liquid state, only 1% NO is non-dimerized 2, and at 100 ° C remains as a dimer 10% N 2 O 4 . (or N 2 O 4 ) reacts in warm water to form nitric acid: 3NO 2 + H 2 O \u003d 2HNO 3 + NO. NO 2 technology therefore it is very essential as an intermediate stage in obtaining an industrially important product– nitric acid.Nitric oxide (V) N 2 O 5 (outdated. nitric anhydride) white crystalline substance, obtained by dehydration of nitric acid in the presence of phosphorus oxide P 4-10: N 2 O 5 dissolves easily in the moisture of the air, re-forming HNO3. Properties of N 2 O 5 determined by the balanceN 2 O 5 is a good oxidizing agent, it reacts easily, sometimes violently, with metals and organic compounds and explodes when heated in its pure state. Likely structure. Evaporation of the solution produces a white explosive with the proposed structure HON=NOH. Nitrous acid HNO 2 is not exists in its pure form, however, aqueous solutions of its low concentration are formed by adding sulfuric acid to barium nitrite:Nitrous acid is also formed by dissolving an equimolar mixture of NO and NO 2 (or N 2 O 3 ) in water. Nitrous acid is slightly stronger than acetic acid. The oxidation state of nitrogen in it is +3 (its structure HON=O), those. it can be both an oxidizing agent and a reducing agent. Under the action of reducing agents, it is usually reduced to NO , and when interacting with oxidizing agents, it is oxidized to nitric acid.

    The rate of dissolution of certain substances, such as metals or iodide ion, in nitric acid depends on the concentration of nitrous acid present as an impurity. Salts of nitrous acid nitrites dissolve well in water, except for silver nitrite.

     NaNO 2 used in the manufacture of dyes.Nitric acid HNO3 one of the most important inorganic products of the mainstream chemical industry. It is used in the technology of many other inorganic and organic substances, such as explosives, fertilizers, polymers and fibers, dyes, pharmaceuticals, etc. see also CHEMICAL ELEMENTS.LITERATURE Azotchik's Handbook. M., 1969
    Nekrasov B.V. Fundamentals of General Chemistry. M., 1973
    Problems of nitrogen fixation. Inorganic and physical chemistry. M., 1982

    Compounds with an oxidation state of –3. Nitrogen compounds in the -3 oxidation state are represented by ammonia and metal nitrides.

    Ammonia- NH 3 is a colorless gas with a characteristic pungent odor. The ammonia molecule has the geometry of a trigonal pyramid with a nitrogen atom at the top. The atomic orbitals of nitrogen are in sp 3- hybrid state. Three orbitals are involved in the formation of nitrogen-hydrogen bonds, and the fourth orbital contains an unshared electron pair, the molecule has a pyramidal shape. The repulsive action of the lone pair of electrons causes the bond angle to decrease from the expected 109.5 to 107.3°.

    At a temperature of -33.4 °C, ammonia condenses to form a liquid with a very high heat of vaporization, which allows it to be used as a refrigerant in industrial refrigeration systems.

    The presence of an unshared electron pair at the nitrogen atom allows it to form another covalent bond according to the donor-acceptor mechanism. Thus, in an acidic environment, the formation of the molecular ammonium cation - NH 4 + occurs. The formation of a fourth covalent bond leads to alignment of bond angles (109.5°) due to the uniform repulsion of hydrogen atoms.

    Liquid ammonia is a good self-ionizing solvent:

    2NH 3 NH 4 + + NH 2 -

    amide anion

    It dissolves alkali and alkaline earth metals, forming colored conductive solutions. In the presence of a catalyst (FeCl 3), the dissolved metal reacts with ammonia to release hydrogen and form an amide, for example:

    2Na + 2NH 3 \u003d 2NaNH 2 + H 2

    sodium amide

    Ammonia is very soluble in water (at 20 °C, about 700 volumes of ammonia dissolve in one volume of water). In aqueous solutions, it exhibits the properties of a weak base.

    NH 3 + H 2 O ® NH 3 × H 2 O NH 4 + + OH -

    = 1.85 10 -5

    In an oxygen atmosphere, ammonia burns with the formation of nitrogen; on a platinum catalyst, ammonia is oxidized to nitric oxide (II):

    4NH 3 + 3O 2 = 2N 2 + 6H 2 O; 4NH 3 + 5O 2 \u003d 4NO + 6H 2 O

    As a base, ammonia reacts with acids to form salts of the ammonium cation, for example:

    NH 3 + HCl = NH 4 Cl

    Ammonium salts are highly soluble in water and slightly hydrolyzed. In the crystalline state, they are thermally unstable. The composition of thermolysis products depends on the properties of the acid forming the salt:

    NH 4 Cl ® NH 3 + HCl; (NH 4) 2 SO 4 ® NH 3 + (NH 4) HSO 4

    (NH 4) 2 Cr 2 O 7 ® N 2 + Cr 2 O 3 + 4H 2 O

    Under the action of alkalis on aqueous solutions of ammonium salts, ammonia is released during heating, which makes it possible to use this reaction as a qualitative one for ammonium salts and as a laboratory method for obtaining ammonia.

    NH 4 Cl + NaOH \u003d NaCl + NH 3 + H 2 O

    In industry, ammonia is obtained by direct synthesis.

    N 2 + 3H 2 2NH 3

    Since the reaction is highly reversible, the synthesis is carried out at elevated pressure (up to 100 MPa). To speed up the process, the process is carried out in the presence of a catalyst (spongy iron promoted by additives) and at a temperature of about 500°C.

    Nitride are formed as a result of the reactions of many metals and non-metals with nitrogen. The properties of nitrides naturally change in a period. For example, for elements of the third period:

    Nitrides of s-elements of groups I and II are crystalline salt-like substances that are easily decomposed by water to form ammonia.

    Li 3 N + 3H 2 O \u003d 3LiOH + NH 3

    Of the halogen nitrides in the free state, only Cl 3 N was isolated, the acid character manifests itself in the reaction with water:

    Cl 3 N + 3H 2 O \u003d 3HClO + NH 3

    The interaction of nitrides of different nature leads to the formation of mixed nitrides:

    Li 3 N + AlN \u003d Li 3 AlN 2; 5Li 3 N + Ge 3 N 4 = 3Li 5 GeN 3

    lithium nitridoaluminate nitridogermanate(IV) lithium

    BN, AlN, Si 3 N 4, Ge 3 N 4 nitrides are solid polymeric substances with high melting points (2000-3000 ° C), they are semiconductors or dielectrics. Nitrides of d-metals - crystalline compounds of variable composition (bertolides), very hard, refractory and chemically stable, exhibit metallic properties: metallic luster, electrical conductivity.

    Compounds with an oxidation state of –2. Hydrazine - N 2 H 4 - the most important inorganic nitrogen compound in the -2 oxidation state.

    Hydrazine is a colorless liquid with a boiling point of 113.5 °C, fuming in air. Hydrazine vapors are extremely toxic and form explosive mixtures with air. Hydrazine is obtained by oxidizing ammonia with sodium hypochlorite:

    2N -3 H 3 + NaCl +1 O \u003d N 2 -2 H 4 + NaCl -1 + H 2 O

    Hydrazine is miscible with water in any ratio and behaves in solution as a weak diacid base, forming two series of salts.

    N 2 H 4 + H 2 O N 2 H 5 + + OH - , K b = 9.3×10 -7 ;

    hydrosonium cation

    N 2 H 5 + + H 2 O N 2 H 6 2+ + OH - , K b = 8.5×10 -15;

    dihydrosonium cation

    N 2 H 4 + HCl N 2 H 5 Cl; N 2 H 5 Cl + HCl N 2 H 6 Cl 2

    hydrosonium chloride dihydrosonium dichloride

    Hydrazine is the strongest reducing agent:

    4KMn +7 O 4 + 5N 2 -2 H 4 + 6H 2 SO 4 \u003d 5N 2 0 + 4Mn +2 SO 4 + 2K 2 SO 4 + 16H 2 O

    Unsymmetrical dimethylhydrazine (heptyl) is widely used as a rocket fuel.

    Compounds with an oxidation state of –1. Hydroxylamine - NH 2 OH - the main inorganic nitrogen compound in the oxidation state -1.

    Hydroxylamine is obtained by reducing nitric acid with hydrogen at the time of isolation during electrolysis:

    HNO 3 + 6H \u003d NH 2 OH + 2H 2 O

    This is a colorless crystalline substance (mp. 33 ° C), highly soluble in water, in which it exhibits the properties of a weak base. With acids it gives hydroxylammonium salts - stable colorless substances soluble in water.

    NH 2 OH + H 2 O + + OH - , K b = 2×10 -8

    hydroxylammonium ion

    The nitrogen atom in the NH 2 OH molecule exhibits an intermediate oxidation state (between -3 and +5), so hydroxylamine can act both as a reducing agent and as an oxidizing agent:

    2N -1 H 2 OH + I 2 + 2KOH = N 0 2 + 2KI + 4H 2 O;

    reducing agent

    2N -1 H 2 OH + 4FeSO 4 + 3H 2 SO 4 = 2Fe 2 (SO 4) 3 + (N -3 H 4) 2 SO 4 + 2H 2 O

    oxidizer

    NH 2 OH easily decomposes when heated, undergoing disproportionation:

    3N -1 H 2 OH \u003d N 0 2 + N -3 H 3 + 3H 2 O;

    Compounds with an oxidation state of +1. Nitric oxide (I) - N 2 O (nitrous oxide, laughing gas). The structure of its molecule can be conveyed by the resonance of two valence schemes, which show that this compound can be considered as nitric oxide (I) only formally, in reality it is nitrogen (V) oxynitride - ON +5 N -3.

    N 2 O is a colorless gas with a slight pleasant smell. In small concentrations it causes bouts of unbridled joy, in large doses it has a general anesthetic effect. A mixture of nitrous oxide (80%) and oxygen (20%) was used in medicine for anesthesia.

    Under laboratory conditions, nitric oxide (I) can be obtained by decomposition of ammonium nitrate. N 2 O obtained by this method contains impurities of higher nitrogen oxides, which are extremely toxic!

    NH 4 NO 3 ¾® N 2 O + 2H 2 O

    According to its chemical properties, nitric oxide (I) is a typical non-salt-forming oxide; it does not react with water, acids and alkalis. When heated, it decomposes to form oxygen and nitrogen. For this reason, N 2 O can act as an oxidizing agent, for example:

    N 2 O + H 2 \u003d N 2 + H 2 O

    Compounds with an oxidation state of +2. Nitric oxide (II) - NO - colorless gas, extremely toxic. In air, it is rapidly oxidized by oxygen to form no less toxic nitric oxide (IV). In industry, NO is produced by the oxidation of ammonia on a platinum catalyst or by passing air through an electric arc (3000-4000 °C).

    4NH 3 + 5O 2 \u003d 4NO + 6H 2 O; N 2 + O 2 \u003d 2NO

    A laboratory method for obtaining nitric oxide (II) is the interaction of copper with dilute nitric acid.

    3Cu + 8HNO 3 (diff.) = 3Cu(NO 3) 2 + 2NO + 4H 2 O

    Nitric oxide (II) is a non-salt-forming oxide, a strong reducing agent, easily reacts with oxygen and halogens.

    2NO + O 2 \u003d 2NO 2; 2NO + Cl 2 = 2NOCl

    nitrosyl chloride

    At the same time, when interacting with strong reducing agents, NO acts as an oxidizing agent:

    2NO + 2H 2 = N 2 + 2H 2 O; 10NO + 4Р = 5N 2 + 2Р 2 O 5

    Compounds with an oxidation state of +3. Nitric oxide (III) - N 2 O 3 - an intensely blue liquid (t.cr. -100 ° C). Stable only in liquid and solid state at low temperatures. It appears to exist in two forms:

    Nitric oxide(III) is obtained by co-condensation of NO and NO 2 vapors. Dissociates in liquids and vapors.

    NO 2 + NO N 2 O 3

    The properties are typical acidic oxide. It reacts with water, forming nitrous acid, with alkalis forms salts - nitrites.

    N 2 O 3 + H 2 O \u003d 2HNO 2; N 2 O 3 + 2NaOH \u003d 2NaNO 2 + H 2 O

    Nitrous acid- medium strength acid (K a = 1×10 -4). It has not been isolated in its pure form, in solutions it exists in two tautomeric forms (tautomers are isomers that are in dynamic equilibrium).

    nitrite form nitro form

    Salts of nitrous acid are stable. The nitrite anion exhibits a pronounced redox duality. Depending on the conditions, it can perform both the function of an oxidizing agent and the function of a reducing agent, for example:

    2NaNO 2 + 2KI + 2H 2 SO 4 = I 2 + 2NO + K 2 SO 4 + Na 2 SO 4 + 2H 2 O

    oxidizer

    KMnO 4 + 5NaNO 2 + 3H 2 SO 4 = 2MnSO 4 + 5NaNO 3 + K 2 SO 4 + 3H 2 O

    reducing agent

    Nitrous acid and nitrites are prone to disproportionation:

    3HN +3 O 2 \u003d HN +5 O 3 + 2N +2 O + H 2 O

    Compounds with an oxidation state of +4. Nitric oxide (IV) - NO 2 - brown gas, with a sharp unpleasant odor. Extremely toxic! In industry, NO 2 is produced by the oxidation of NO. The laboratory method for obtaining NO 2 is the interaction of copper with concentrated nitric acid, as well as the thermal decomposition of lead nitrate.

    Cu + 4HNO 3 (conc.) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O;

    2Pb(NO 3) 2 \u003d 2PbO + 4NO 2 + O 2

    The NO 2 molecule has one unpaired electron and is a stable free radical, so nitric oxide easily dimerizes.

    The dimerization process is reversible and very sensitive to temperature:

    paramagnetic, diamagnetic,

    brown colorless

    Nitrogen dioxide is an acidic oxide that reacts with water to form a mixture of nitric and nitrous acid (mixed anhydride).

    2NO 2 + H 2 O \u003d HNO 2 + HNO 3; 2NO 2 + 2NaOH \u003d NaNO 3 + NaNO 2 + H 2 O

    Compounds with an oxidation state of +5. Nitric oxide (V) - N 2 O 5 - white crystalline substance. Obtained by dehydration of nitric acid or oxidation of nitric oxide (IV) with ozone:

    2HNO 3 + P 2 O 5 \u003d N 2 O 5 + 2HPO 3; 2NO 2 + O 3 \u003d N 2 O 5 + O 2

    In the crystalline state, N 2 O 5 has a salt-like structure - + -, in vapor (t. vozg. 33 ° C) - molecular.

    N 2 O 5 - acid oxide - nitric acid anhydride:

    N 2 O 5 + H 2 O \u003d 2HNO 3

    Nitric acid- HNO 3 - a colorless liquid with a boiling point of 84.1 ° C, decomposes when heated and in the light.

    4HNO 3 \u003d 4NO 2 + O 2 + 2H 2 O

    Nitrogen dioxide impurities give concentrated nitric acid a yellow-brown color. Nitric acid is miscible with water in any ratio and is one of the strongest mineral acids; it completely dissociates in solution.

    The structure of the nitric acid molecule is described by the following structural formulas:

    Difficulties with writing the structural formula of HNO 3 are caused by the fact that, showing in this compound the oxidation state +5, nitrogen, as an element of the second period, can form only four covalent bonds.

    Nitric acid is one of the strongest oxidizing agents. The depth of its recovery depends on many factors: concentration, temperature, reducing agent. Usually, when oxidized with nitric acid, a mixture of reduction products is formed:

    HN +5 O 3 ® N +4 O 2 ® N +2 O ® N +1 2 O ® N 0 2 ® +

    The predominant product of the oxidation of non-metals and inactive metals with concentrated nitric acid is nitric oxide (IV):

    I 2 + 10HNO 3 (conc) = 2HIO 3 + 10NO 2 + 4H 2 O;

    Pb + 4HNO 3 (conc) = Pb (NO 3) 2 + 2NO 2 + 2H 2 O

    Concentrated nitric acid passivates iron and aluminum. Aluminum is passivated even with dilute nitric acid. Nitric acid of any concentration does not affect gold, platinum, tantalum, rhodium and iridium. Gold and platinum are dissolved in aqua regia - a mixture of concentrated nitric and hydrochloric acids in a ratio of 1: 3.

    Au + HNO 3 + 4HCl \u003d H + NO + 2H 2 O

    The strong oxidizing effect of aqua regia is due to the formation of atomic chlorine during the decomposition of nitrosyl chloride, a product of the interaction of nitric acid with hydrogen chloride.

    HNO 3 + 3HCl \u003d Cl 2 + NOCl + 2H 2 O;

    NOCl = NO + Cl×

    An effective solvent for low-active metals is a mixture of concentrated nitric and hydrofluoric acids.

    3Ta + 5HNO 3 + 21HF = 3H 2 + 5NO + 10H 2 O

    Diluted nitric acid, when interacting with non-metals and low-active metals, is reduced mainly to nitric oxide (II), for example:

    3P + 5HNO 3 (razb) + 2H 2 O \u003d 3H 3 PO 4 + 5NO;

    3Pb + 8HNO 3 (razb) \u003d 3Pb (NO 3) 2 + 2NO + 4H 2 O

    Active metals reduce dilute nitric acid to N 2 O, N 2 or NH 4 NO 3, for example,

    4Zn + 10HNO 3 (razb) \u003d 4Zn (NO 3) 2 + NH 4 NO 3 + 3H 2 O

    The bulk of nitric acid goes to the production of fertilizers and explosives.

    Nitric acid is produced industrially by the contact or arc method, which differ in the first stage - the production of nitric oxide (II). The arc method is based on the production of NO by passing air through an electric arc. In the contact process, NO is produced by the oxidation of ammonia with oxygen over a platinum catalyst. Next, nitric oxide (II) is oxidized to nitric oxide (IV) by atmospheric oxygen. By dissolving NO 2 in water in the presence of oxygen, nitric acid is obtained with a concentration of 60-65%.

    4NO 2 + O 2 + 2H 2 O \u003d 4HNO 3

    If necessary, nitric acid is concentrated by distillation with concentrated sulfuric acid. In the laboratory, 100% nitric acid can be obtained by the action of concentrated sulfuric acid on crystalline sodium nitrate when heated.

    NaNO 3 (cr) + H 2 SO 4 (conc) = HNO 3 + NaHSO 4

    Salts of nitric acid- nitrates - highly soluble in water, thermally unstable. The decomposition of nitrates of active metals (excluding lithium), which are in the series of standard electrode potentials to the left of magnesium, leads to the formation of nitrites. For example:

    2KNO 3 \u003d 2KNO 2 + O 2

    During the decomposition of lithium, magnesium nitrates, as well as metal nitrates, located in the series of standard electrode potentials to the right of magnesium, up to copper, a mixture of nitric oxide (IV) and oxygen is released. For example:

    2Cu(NO 3) 2 \u003d 2CuO + 4NO 2 + O 2

    Nitrates of metals located at the end of the activity series decompose to free metal:

    2AgNO 3 \u003d 2Ag + 2NO 2 + O 2

    Sodium, potassium and ammonium nitrates are widely used for the production of gunpowder and explosives, as well as nitrogen fertilizers (saltpeter). Ammonium sulfate, ammonia water and carbamide (urea) - full carbonic acid amide are also used as fertilizers:

    Hydrogen azide(dinitridonitrate) - HN 3 (HNN 2) - a colorless volatile liquid (mp -80 ° C, bp 37 ° C) with a pungent odor. The central nitrogen atom is in sp hybridization, the oxidation state is +5, the atoms adjacent to it have an oxidation state of –3. Molecule structure:

    An aqueous solution of HN 3 - hydronitrous acid is close in strength to acetic acid, K a \u003d 2.6 × 10 -5. Stable in dilute solutions. It is obtained by the interaction of hydrazine and nitrous acid:

    N 2 H 4 + HNO 2 \u003d HN 3 + 2H 2 O

    In terms of oxidizing properties, HN 3 (HN +5 N 2) resembles nitric acid. So, if the interaction of a metal with nitric acid produces nitric oxide (II) and water, then with hydrazoic acid - nitrogen and ammonia. For example,

    Cu + 3HN +5 N 2 = Cu(N 3) 2 + N 2 0 + NH 3

    A mixture of HN 3 and HCl behaves like aqua regia. Salts of hydronitrous acid - azides. Only alkali metal azides are relatively stable; at temperatures > 300 °C they are destroyed without explosion. The rest disintegrate with an explosion on impact or heating. Lead azide is used in the production of detonators:

    Pb(N 3) 2 = Pb + 3N 2 0

    The starting product for the production of azides is NaN 3, which is formed as a result of the reaction of sodium amide and nitric oxide (I):

    NaNH 2 + N 2 O \u003d NaN 3 + H 2 O

    4.2 Phosphorus

    Phosphorus is represented in nature by one isotope - 31 P, the clarke of phosphorus is 0.05 mol.%. It occurs in the form of phosphate minerals: Ca 3 (PO 4) 2 - phosphorite, Ca 5 (PO 4) 3 X (X \u003d F, Cl, OH) - apatites. It is part of the bones and teeth of animals and humans, as well as the composition of nucleic acids (DNA and RNA) and adenosine phosphoric acids (ATP, ADP and AMP).

    Phosphorus is obtained by reduction of phosphorite with coke in the presence of silicon dioxide.

    Ca 3 (PO 4) 2 + 3SiO 2 + 5C = 3CaSiO 3 + 2P + 5CO

    A simple substance - phosphorus - forms several allotropic modifications, of which the main ones are white, red and black phosphorus. White phosphorus is formed during the condensation of phosphorus vapor and is a white wax-like substance (mp 44 ° C), insoluble in water, soluble in some organic solvents. White phosphorus has a molecular structure and consists of tetrahedral molecules P 4 .

    The bond strength (valence angle P-P-P is only 60 °) determines the high reactivity and toxicity of white phosphorus (lethal dose is about 0.1 g). Since white phosphorus is highly soluble in fats, milk cannot be used as an antidote for poisoning. In air, white phosphorus spontaneously ignites, so it is stored in a hermetically sealed chemical container under a layer of water.

    Red phosphorus has a polymer structure. It is obtained by heating white phosphorus or irradiating it with light. Unlike white phosphorus, it is slightly reactive and non-toxic. However, residual amounts of white phosphorus can make red phosphorus toxic!

    Black phosphorus is obtained by heating white phosphorus under a pressure of 120 thousand atm. It has a polymer structure, has semiconductor properties, is chemically stable and non-toxic.

    Chemical properties. White phosphorus is spontaneously oxidized by atmospheric oxygen at room temperature (oxidation of red and black phosphorus occurs when heated). The reaction proceeds in two stages and is accompanied by luminescence (chemiluminescence).

    2P + 3O 2 \u003d 2P 2 O 3; P 2 O 3 + O 2 \u003d P 2 O 5

    Phosphorus also reacts stepwise with sulfur and halogens.

    2P + 3Cl 2 \u003d 2PCl 3; PCl 3 + Cl 2 = PCl 5

    When interacting with active metals, phosphorus acts as an oxidizing agent, forming phosphides - phosphorus compounds in the -3 oxidation state.

    3Ca + 2P = Ca 3 P 2

    Oxidizing acids (nitric and concentrated sulfuric acids) oxidize phosphorus to phosphoric acid.

    P + 5HNO 3 (conc) = H 3 PO 4 + 5NO 2 + H 2 O

    When boiling with alkali solutions, white phosphorus disproportionates:

    4P 0 + 3KOH + 3H 2 O = P -3 H 3 + 3KH 2 P +1 O 2

    phosphine potassium hypophosphite

    Nitrogen is perhaps the most common chemical element in the entire solar system. To be more specific, nitrogen is the 4th most abundant. Nitrogen in nature is an inert gas.

    This gas is colorless and odorless and very difficult to dissolve in water. However, nitrate salts tend to react very well with water. Nitrogen has a low density.

    Nitrogen is an amazing element. There is an assumption that it got its name from the ancient Greek language, which means “lifeless, spoiled” in translation from it. Why such a negative attitude towards nitrogen? After all, we know that it is part of proteins, and breathing without it is almost impossible. Nitrogen plays an important role in nature. But in the atmosphere this gas is inert. If it is taken as it is in its original form, then many side effects are possible. The victim may even die from suffocation. After all, nitrogen is called lifeless because it does not support combustion or respiration.

    Under normal conditions, such a gas reacts only with lithium, forming a compound such as lithium nitride Li3N. As we can see, the oxidation state of nitrogen in such a compound is -3. With other metals, and of course, it also reacts, but only when heated or when using various catalysts. By the way, -3 is the lowest oxidation state of nitrogen, since only 3 electrons are needed to completely fill the outer energy level.

    This indicator has various meanings. Each oxidation state of nitrogen has its own compound. It is better to just remember such connections.

    5 - the highest degree of oxidation of nitrogen. Occurs in and in all nitrate salts.



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