Oxygen-containing nitrogen acids. Nitrogen and its compounds

Acid (H n A)

Acid residue (A)

Corresponding acid oxide

H 2 SO 4 sulfuric

SO 4 (II) sulfate

SO3 sulfur oxide (VI)

HNO 3 nitrogen

NO3(I)nitrate

N 2 O 5 nitric oxide (V)

HMnO 4 manganese

MnO 4 (I) permanganate

Mn2O7 manganese oxide ( VII)

H 2 SO 3 sulfurous

SO 3 (II) sulfite

SO2 sulfur oxide (IV)

H 3 PO 4 orthophosphoric

PO 4 (III) orthophosphate

P 2 O 5 phosphorus oxide (V)

HNO 2 nitrogenous

NO 2 (I) nitrite

N 2 O 3 nitric oxide (III)

H 2 CO 3 coal

CO 3 (II) carbonate

CO2 carbon monoxide ( IV)

H 2 SiO 3 silicon

SiO 3 (II) silicate

SiO 2 silicon(IV) oxide

HClO hypochlorous

ClO(I) hypochlorite

C l 2 O chlorine oxide (I)

HClO 2 chloride

ClO 2 (I) chlorite

C l 2 O 3 chlorine oxide (III)

HClO 3 chlorate

ClO 3 (I) chlorate

C l 2 O 5 chlorine oxide (V)

HClO 4 chlorine

ClO 4 (I) perchlorate

C l 2 O 7 chlorine oxide (VII)

Acid (H n A)

Acid residue (A)

HCl hydrochloric, hydrochloric

Cl(I) chloride

H 2 S hydrogen sulfide

S(II) sulfide

HBr hydrogen bromide

Br(I) bromide

HI hydrogen iodide

I(I)iodide

HF hydrogen fluoride, fluoride

F(I) fluoride

oxygen-free

oxygen-containing

HCl, HBr, HI, HF, H2S

HNO 3, H 2 SO 4 and others

RECEIVING

1. Direct interaction of nonmetals

H 2 + Cl 2 = 2 HCl

1. Acidic oxide + water = acid

SO 3 + H 2 O = H 2 SO 4

2. Exchange reaction between salt and less volatile acid

2 NaCl (tv.) + H 2 SO 4 (conc.) = Na 2 SO 4 + 2HCl

Indicator name

Neutral environment

Acidic environment

Litmus

Violet

Red

Phenolphthalein

Colorless

Colorless

Methyl orange

Orange

Red

Universal indicator paper

Orange

Red

N 2 + 3H 2 ↔2NH 3

To obtain nitrogen oxide (II) from ammonia, the latter is oxidized with oxygen:

4NH 3 + 5O 2 →4NO + 6H 2 O

Nitric oxide (IV) is obtained from nitric oxide (II) by its oxidation with oxygen:

2NO + O 2 →2NO 2

When nitric oxide (IV) reacts with water in the presence of oxygen, nitric acid is obtained:

4NO 2 +2H 2 O + O 2 →4HNO 3

When nitric acid reacts with ammonia solution, ammonium nitrate is obtained:

HNO 3 + NH 3 →NH 4 NO 3

When heated, ammonium nitrate breaks down into nitric oxide (I) and water.

Oxygen-containing acids are also classified as hydroxides. These are electrolytes that form when dissociated into aqueous solutions of the positively charged ions, only hydrogen ions H +, or, more precisely, hydronium ions H 3 O + - hydrated hydrogen ion. More general definition: acids– these are substances that donor H + protons. Depending on the number of hydrogen cations formed during the dissociation of the acid, acids are also classified as bases, according to their basicity. There are one-, two-, three- and tetrabasic acids. For example, nitric acid HNO 3, nitrous acid HNO 2 are monobasic acids, carbonic acid H 2 CO 3, sulfuric acid H 2 SO 4 are dibasic acids, orthophosphoric acid H 3 PO 4 is a tribasic acid, and orthosilicic acid H 4 SiO 4 is tetrabasic acid.

Nomenclature of oxygen-containing acids: By international systematic nomenclature the names of oxygen-containing acids are formed, as indicated earlier, taking into account the anion included in the acid. For example:

H 3 PO 4 - trihydrogen tetraoxophosphate (V) or trihydrogen orthophosphate

H 2 CO 3 - dihydrogen trioxocarbonate (IV)

HNO 3 - hydrogen trioxonitrate (V)

H 2 SiO 3 - dihydrogen trioxosilicate (IV) or dihydrogen metasilicate

H 2 SO 4 - dihydrogen tetraoxosulfate(VI) (the number of hydrogen atoms in acids may not be indicated)

According to systematic nomenclature, the names of acids are rarely used; most often, traditionally established names are used, which are formed from Russian name of the element (Russian nomenclature) according to certain rules (see table). The table shows a list of oxygen-containing acids, the salts of which are most common in nature. Please note that the title acid residue determines the name of the salt and builds it most often according to semi-systematic (international) nomenclature from the Latin name of the element. In this regard, it is necessary to remember the Latin names of the elements most often found in acids, for example, N - nitrogen, in Russian transcription of the Latin name it sounds like [nitrogenium], C - carbon - [carbonium], S - sulfur - [sulfur], Si- silicon - [silicium], tin - [stannum], lead - [plumbum], arsenic - [arsenicum], etc. The table shows general rules according to which most inorganic oxygen-containing acids of other elements, their acid residues and salts can be named.

Table of the most common oxygen-containing acids

Hydrosols and the names of their acidic residues will be discussed in the “salts” section. The rules for naming oxygen-containing acids and acid residues (except for those that have trivial names or should be called by systematic nomenclature) are as follows:

highest s. O. element (equal to group number in periodic table) – root of the Russian name of the element + ending “ A I" or "s A I"

N

With. O. – oxidation state

oxygen-containing

acids

S.o. element< max – корень русского названия элемента +

ending " And flock" or "s And flock"

highest s.o. element – ​​root of the Latin name of the element +

Name suffix " A T"

acidic

remainder

s.o. element< max – Latin name element + suffix " And T"

Knowing the above rules, it is easy to derive acid formulas for various elements (taking into account their position in the periodic table) and name them. For example, the metal Sn is tin (1V gr.), the Latin name is stannum (“stannum”):

Max s.o. = +4 Min s.o. = +2

Oxides: SnO 2 SnO

amphote. amphote.

+H 2 ABOUT +H 2 ABOUT

H 2 SnO 3 H 2 SnO 2

tin oh tin acid exhausted acid

SnO 3 2- SnO 2 2-

Stann at- ion, stann it-ion,

Na 2 SnO 3 – stannate Na Na 2 SnO 2 – stannit Na

The oxides of some elements correspond to two acids: meta- And orthoacid, formally they differ by one H2O molecule.

Derivation of the formula meta and ortho acids(if they exist for a given element): with the formal addition of one H 2 O molecule to the oxide, we obtain the formula of a meta-acid, the subsequent addition of another water molecule to the formula of the meta-acid allows us to derive the formula of an ortho acid. For example, let’s derive the formula for meta- and ortho acids corresponding to P(V) oxide:

+ H 2 O + H 2 O

H 2 P 2 O 6  HPO 3 - metaphosphoric acid H 3 PO 4 - orthophosphoric acid

Let's give an example of an inverse problem: name the salts NaBO 2 and K 3 BO 3. The oxidation state of the boron atom in these salts is +3 (check the calculation), therefore, the salts are formed from the acidic oxide B 2 O 3. If in both salts the oxidation states of boron are the same, but the types of acidic residues are different, then these are salts of meta- and orthoboric acid. Let us derive the formulas of these acids:

B 2 O 3 HBO 2

+ N 2 ABOUT + N 2 ABOUT

HBO 2 - metaboric acid, H 3 BO 3 - orthoboric acid,

salts – salt metaborates – orthoborates

Names of salts: NaBO 2 – sodium metaborate; Na 3 BO 3 - sodium orthoborate.

Structure and chemical properties oxygen-containing nitrogen compounds.

Nitric oxide (I) N2 O-laughing gas. Colorless gas with a sweetish taste. Does not react chemically with water. Chemically inactive. It does not react with water, acids, alkalis, halogens or ozone.

At elevated temperatures it decomposes: 2N2O = 2 N2 + O2 At elevated temperatures, it is a strong oxidizing agent.

Oxidizes Me P, C, S. Ni+N2O=NiO+N2 .N2O + Cu = CuO + N2

When a mixture of nitrogen (I) oxide and ammonia is ignited, an explosion occurs: 3N2O+2NH3=4N2+3H2O

2NO+O2=2NO2 - second stage of nitric acid

When interacting with strong oxidizing agents, N2O can exhibit the properties of a reducing agent: 5N2O+8KMnO4+7H2SO4=5Mn(NO3)2+3MnSO4+4K2SO4+7H2O

Nitric oxide (II) NO- colorless gas is a typical reducing agent. The only gas that can be formed at 3000C is N2+O2=2NO. It does not form acid. Does not react with water

At temperatures above 1000 C, it decomposes: NO = N2 + O2

NO is also characterized by addition reactions; in this reaction, NO exhibits the properties of a reducing agent with the formation of nitrosyl chloride: 2NO+Cl2=2NOCl

In the presence of stronger reducing agents, NO exhibits oxidizing properties: 2NO+2H2S=N2+2S+2H2O 2NO+2CO=N2+CO2

At the same time, a mixture of equal volumes of NO and H2 explodes when heated: 2NO + 2H2 = N2 + 2H2O

Nitric oxide(3)N2 O3 unstable, exists only at low temperatures. Bright blue. At 0C it decomposes: N2O3=NO+NO2 N2O3+H2O=2HNO3 N2O3+2KOH=2KNO2+H2O

nitric oxide(4)NO2 - brown gas or N2O4 - colorless. NO2 (drill when heating) = N2O4 (when cooling)

Reacts with water: 3NO2 + H2O = 2HNO3 + NO, Na2O4+H2O=HNO3+HNO2.

4NO2+2H2O+O2=4HNO3-3 stage of obtaining nitric acid exhibits the properties of a reducing agent

When NO2 is dissolved in alkalis, both nitrates and nitrites are formed: 2NO2+2KOH=KNO3+KNO2+H2O

Liquid NO2 is used to obtain anhydrous nitrates: Zn+2N2O4=Zn(NO3)2+2NO

interacts with non-metals (phosphorus, sulfur and carbon burn in it). In these reactions NO2 is an oxidizing agent: 2NO2+C=CO2+2NO, 2NO2+4HCl=NOCl+H2O+Cl2

Nitric oxide (V) N2 O5 volatile, hygroscopic, colorless, unstable.

Already at room temperature it gradually decomposes: N2O5=NO2+O2

Very strong oxidizing agent: N2O5+I2=I2O5+N2. Many organic matter upon contact with it, they ignite.

When dissolved in water, nitric acid is formed: N2O5+H2O=HNO3

Dissolves in alkalis to form nitrates: N2O5+2NaOH=2NaNO3+H2O

Oxoacids:

Nitrous acidHNO2

It is a weak acid and is known only in highly dilute aqueous solutions.

2 HNO2+ 2 HI = I2+ 2 NO + 2 H2O

3 HNO2↔HNO3+ 2 NO + H2O

When the solution is concentrated or heated, it decomposes: 2HNO2=NO+NO2+H2O

Exhibits redox duality. Under the influence of reducing agents it is reduced, and in reactions with oxidizing agents it is oxidized to HNO3: HNO2+Cl2+H2O=HNO3+2HCl
2HNO2+2KI+H2SO4=2NO+I2+K2SO4+2H2O

5HNO2+2KMnO4+3H2SO4=5HNO3+2MnSO4+K2SO4+3H2O

Prone to disproportionation reactions: 3HNO2=HNO3+2NO+H2O

Nitric acidHNO3 already decomposes under the influence of light: HNO3=4NO2+O2+2H2O

is one of the strongest acids.

Nitrogen acts on almost all metals (except gold, platinum, tantalum, rhodium, iridium), turning them into nitrates, and some metals into oxides.

Cu+HNO3(conc)=Cu(NO3)2+NO2+H2O. Cu+HNO3(dil)=Cu(NO3)2+NO+H2O

Mg+HNO3(dil)=Mg(NO3)2+N2O+H2O. Zn+HNO3(very dilute)=Zn(NO3)2+NH4NO3+H2O

Nitrogen oxides

Nitrous acid, its salts

Nitric acid and its salts

1. Nitrogen oxides

acid nitrous oxide salt

There are five nitrogen oxides in total: N2O, NO – non-salt-forming oxides; N2O3, NO2, N2O5 – acidic oxides.

A) N2 O– nitrous oxide.

Obtained by the decomposition of ammonium nitrate at 250 o C.

NH4NO3 → N2O + 2H2O

It is a gas with a pleasant faint odor. Inhaling small amounts of this gas has an intoxicating effect, hence the name “laughing gas.” In large doses it causes loss of pain sensitivity. The N2O molecule has a linear structure

N2O dissolves well in H2O, but does not form stable compounds. It does not react with water, acids or alkalis. Even with low heating, N2O decomposes, releasing O2.

2N2O → 2N2 + O2

Therefore, N2O is an oxidizing agent in relation to all substances that directly react with oxygen.

N2O + H2 = N2 + H2O.

b) NO– nitric oxide(P). Also a non-salt-forming oxide. Under normal conditions, NO is a colorless gas. In industry, they are obtained by the oxidation of H3N on a platinum catalyst when heated:

4NH3 + 5O2 = 4NO + 6H2O

In the laboratory, NO is produced by the action of dilute HNO3 on Cu:

3Cu + 8HNO3dil. = 3Cu(NO3)2 + 2NO + 4H2O.

Unlike all other nitrogen oxides, NO is also formed by direct interaction of simple substances:

The structural formula of NO is as follows: one electron in NO is antibonding, and 6 electrons are bonding, that is, the bond order is 2.5. The NO molecule is quite stable and its decomposition is noticeable only at 500°C. NO is a chemically active compound with redox duality. Under the influence of O2 in air, it is easily oxidized to NO2, and is also oxidized by halogens:

2NO + O2 = 2NO2; 2NO + Cl2 = 2NOCl

As an oxidizing agent, NO easily oxidizes SO2 to SO3;

2SO2 + 2NO = 2SO3 + N2

With hydrogen (especially when equal volumes) NO explodes when heated:

2NO + 2H2 = N2 + 2H2O

NO is slightly soluble in water and does not react with water.

V) NO2 – nitric oxide (IV) - red-brown poisonous gas with a characteristic odor.

Its molecule has an angular shape, the bond order between N and O is 1.5. The NO2 molecule, even in vapor, is partially dimerized:

2NO2 ↔ N2O4 + Q

These two compounds are in equilibrium with each other at temperatures from –11o to 140oC. The NO2 molecule is characterized by high chemical activity. As the temperature rises, NO2 is one of the most energetic oxidizing agents (C, S, P burn in it). At temperatures above 500°C, NO2 decomposes into NO and O2. When dissolved in water, two acids are formed:

2N+4O2 + H2O → HN+5O3 + HN+3O2,

that is, NO2 is a mixed anhydride of nitric and nitrous acids.

Similarly with alkali:

2NO2 + 2NaOH = NaNO3 + NaNO2 + H2O

4NO2 + O2 + 2H2O = 4HNO3

(this reaction is used industrially to produce HNO3).

In laboratory conditions, NO2 is obtained:

or thermal decomposition of nitrates:

2Рв(NO3)2 = 2РвО + 4NO2 + O2

G). N2 O5 – nitric oxide (V) – Nitric anhydride is obtained by dehydration of nitric acid with phosphoric anhydride (careful dehydration) or oxidation of NO2 with ozone.

2HNO3 + P2O5 = 2HPO3 + N2O52NO2 + O3 → N2O5 + O2.

N2O5 is a white crystalline substance. At room temperature, N2O5 gradually decomposes into NO2 and O2, and explodes when heated:

2N2O5 = 4NO2 + O2

When interacting with H2O, it forms nitric acid:

N2O5 + H2O → 2HNO3

N2O5 is a very strong oxidizing agent. Many organic substances ignite upon contact with it.

d). N2 O3 – nitric oxide (N)– nitrogenous anhydride, formed by the reaction:

NO2 + NO ↔ N2O3.

The equilibrium of this reaction, even at 25°C, is shifted to the left, that is, N2O3 is an unstable compound. Exists only at low temperatures in the solid state (light blue crystals). In the form of liquid and vapor it strongly dissociates:

N2O3 ↔ NO2 + NO

N2O3 can also be obtained from the decomposition of HNO2, which is very unstable:

2HNO2 ↔ H2O + N2O3

N2O3 is an acidic oxide, therefore it easily reacts with alkalis:

N2O3 + 2NaOH = 2NaNO2 + H2O

When dissolved in water, HNO2 is obtained:

N2O3 + HOH ↔ 2HNO2

Structure of N2O3:

2 . Nitrous acid, ee salt

Nitrous acid HNO2 is known only in dilute aqueous solutions. IN pure form does not exist. When heated, it decomposes: 2HNO2 = NO + NO2 + H2O.

HNO2 is an acid of medium strength (K ≈ 5∙10-4).

The HNO2 molecule exists in two tautomeric forms:

Nitrites of metal elements are quite stable, and nitrites of alkali metals even sublime without decomposition.

Nitrogen in HNO2 has C.O. = +3, that is, an intermediate oxidation state, therefore both the acid and the salt have redox duality. Strong oxidizing agents convert NO2- to NO3-:

5NaNO2 + 2KMnO4 + 3H2SO4 → 5NaNO3 + 2MnSO4 + K2SO4 + 3H2O

Strong reducing agents usually reduce NO2- to NO:

2NaNO2 + 2KI + 2H2SO4 → Na2SO4 + 2NO + K2SO4 + I2 + 2H2O

In addition, nitrogen compounds (N) are prone to disproportionation reactions:

3HNO2 = HNO3 + 2NO + H2O

2HNO2 = NO + NO2 + H2O

3. Nitric acid and its salts

Nitric acid HNO3 is produced industrially by the catalytic oxidation of NH3 to NO, then NO is oxidized with atmospheric oxygen to NO2, and then a mixture of NO2 with excess air is absorbed with water (or dilute HNO3).

4NH3 + 5O2 → 4NO + 6H2O

2NO + O2 → 2NO2

4NO2 + O2 + 2H2O → 4HNO3

In the laboratory, HNO3 is obtained by concentrated action. H2SO4 to sodium nitrate:

NaNO3 + H2SO4 = NaHSO4 + HNO3

Under normal conditions, HNO3 is a colorless liquid (ρ = 1.52 g/cm3), boiling at 84.1 °C. HNO3 mixes with water in any ratio. In aqueous solution, HNO3 is a strong acid that is almost completely dissociated. During storage, HNO3 (especially when heated and illuminated) decomposes: 4HNO3 = 4NO2 + O2 + 2H2O.

In the air it “smoke”, as its vapors with air moisture form small droplets of fog.

HNO3 has a flat structure:

The covalency of nitrogen in HNO3 is 4.

HNO3 is a strong oxidizing agent. It destroys animal and plant tissues, its vapors are poisonous. Oxidizes many metals and non-metals:

Cu + 4HNO3 (conc.) = Cu(NO3)2 + 2NO2 + 2H2O

3Cu + 8HNO3 (diluted) = 3Cu(NO3)2 + 2NO + 4H2O

4Zn + 10HNO3 (very dilute) = 4Zn(NO3)2 + NH4NO3 + 3H2O

S + 6HNO3 (conc.) = H2SO4 + 6NO2 + 2H2O

3P + 5HNO3 (dil.) + 2H2O = 3H3PO4 + 5NO

A mixture of one volume of HNO3 and three volumes of concentrated HCl is called “aqua regia.” It is a stronger oxidizing agent than HNO3 and reacts with noble metals such as gold and platinum, converting them into complex chlorides:

Au + HNO3 + 4HCl = NO + 2H2O + H.