Non-salt-forming (indifferent, indifferent) oxides CO, SiO, N 2 0, NO.

Salt-forming oxides:

Basic. Oxides whose hydrates are bases. Metal oxides with oxidation states +1 and +2 (less often +3). Examples: Na 2 O - sodium oxide, CaO - calcium oxide, CuO - copper (II) oxide, CoO - cobalt (II) oxide, Bi 2 O 3 - bismuth (III) oxide, Mn 2 O 3 - manganese (III) oxide ).

Amphoteric. Oxides whose hydrates are amphoteric hydroxides. Metal oxides with oxidation states +3 and +4 (less often +2). Examples: Al 2 O 3 - aluminum oxide, Cr 2 O 3 - chromium (III) oxide, SnO 2 - tin (IV) oxide, MnO 2 - manganese (IV) oxide, ZnO - zinc oxide, BeO - beryllium oxide.

Acidic. Oxides whose hydrates are oxygen-containing acids. Non-metal oxides. Examples: P 2 O 3 - phosphorus oxide (III), CO 2 - carbon oxide (IV), N 2 O 5 - nitrogen oxide (V), SO 3 - sulfur oxide (VI), Cl 2 O 7 - chlorine oxide ( VII). Metal oxides with oxidation states +5, +6 and +7. Examples: Sb 2 O 5 - antimony (V) oxide. CrOz - chromium (VI) oxide, MnOz - manganese (VI) oxide, Mn 2 O 7 - manganese (VII) oxide.

Change in the nature of oxides with increasing oxidation state of the metal

Physical properties

Oxides are solid, liquid and gaseous, of different colors. For example: copper (II) oxide CuO is black, calcium oxide CaO is white - solids. Sulfur oxide (VI) SO 3 is a colorless volatile liquid, and carbon monoxide (IV) CO 2 is a colorless gas under ordinary conditions.

State of aggregation

CaO, CuO, Li 2 O and other basic oxides; ZnO, Al 2 O 3, Cr 2 O 3 and other amphoteric oxides; SiO 2, P 2 O 5, CrO 3 and other acid oxides.

SO 3, Cl 2 O 7, Mn 2 O 7, etc.

Gaseous:

CO 2, SO 2, N 2 O, NO, NO 2, etc.

Solubility in water

Soluble:

a) basic oxides of alkali and alkaline earth metals;

b) almost all acid oxides (exception: SiO 2).

Insoluble:

a) all other basic oxides;

b) all amphoteric oxides

Chemical properties

1. Acid-base properties

Common properties of basic, acidic and amphoteric oxides are acid-base interactions, which are illustrated by the following diagram:

(only for oxides of alkali and alkaline earth metals) (except SiO 2).

Amphoteric oxides, having the properties of both basic and acidic oxides, interact with strong acids and alkalis:

2. Redox properties

If an element has a variable oxidation state (s.o.), then its oxides with low s. O. can exhibit reducing properties, and oxides with high c. O. - oxidative.

Examples of reactions in which oxides act as reducing agents:

Oxidation of oxides with low c. O. to oxides with high c. O. elements.

2C +2 O + O 2 = 2C +4 O 2

2S +4 O 2 + O 2 = 2S +6 O 3

2N +2 O + O 2 = 2N +4 O 2

Carbon (II) monoxide reduces metals from their oxides and hydrogen from water.

C +2 O + FeO = Fe + 2C +4 O 2

C +2 O + H 2 O = H 2 + 2C +4 O 2

Examples of reactions in which oxides act as oxidizing agents:

Reduction of oxides with high o. elements to oxides with low c. O. or to simple substances.

C +4 O 2 + C = 2C +2 O

2S +6 O 3 + H 2 S = 4S +4 O 2 + H 2 O

C +4 O 2 + Mg = C 0 + 2MgO

Cr +3 2 O 3 + 2Al = 2Cr 0 + 2Al 2 O 3

Cu +2 O + H 2 = Cu 0 + H 2 O

The use of oxides of low-active metals for the oxidation of organic substances.

Some oxides in which the element has an intermediate c. o., capable of disproportionation;

For example:

2NO 2 + 2NaOH = NaNO 2 + NaNO 3 + H 2 O

Methods of obtaining

1. Interaction of simple substances - metals and non-metals - with oxygen:

4Li + O 2 = 2Li 2 O;

2Cu + O 2 = 2CuO;

4P + 5O 2 = 2P 2 O 5

2. Dehydration of insoluble bases, amphoteric hydroxides and some acids:

Cu(OH) 2 = CuO + H 2 O

2Al(OH) 3 = Al 2 O 3 + 3H 2 O

H 2 SO 3 = SO 2 + H 2 O

H 2 SiO 3 = SiO 2 + H 2 O

3. Decomposition of some salts:

2Cu(NO 3) 2 = 2CuO + 4NO 2 + O 2

CaCO 3 = CaO + CO 2

(CuOH) 2 CO 3 = 2CuO + CO 2 + H 2 O

4. Oxidation of complex substances with oxygen:

CH 4 + 2O 2 = CO 2 + H 2 O

4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2

4NH 3 + 5O 2 = 4NO + 6H 2 O

5. Reduction of oxidizing acids with metals and non-metals:

Cu + H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O

10HNO 3 (conc) + 4Ca = 4Ca(NO 3) 2 + N 2 O + 5H 2 O

2HNO 3 (diluted) + S = H 2 SO 4 + 2NO

6. Interconversions of oxides during redox reactions (see redox properties of oxides).

Such a weak chemical interaction, which we classify as type VI, can be expressed by the following scheme:

Me"" m ABOUT n= m[Me""] Me" + n[O] Me" ,

where is Me"" m ABOUT n- ceramic or glass oxide; [Me""] Me" and [O] Me" are solid solutions of metal and oxygen that form ceramic oxide in the metal being welded to it, respectively.

An interaction of this type can occur when there is a large difference in the Gibbs energy of the formation of the ceramic or glass oxide and the oxide of the metal being welded.

The possibility of an interaction of this type is indicated, for example, by the phenomena of coagulation of strengthening phases (intermetallic compounds, oxides, carbides, carbonitrides), occurring at elevated temperatures in dispersion-strengthened materials due to the dissolution of small particles in the matrix and the growth of large ones. The possibility and degree of such interaction between the strengthener and the matrix determine the heat resistance of composite materials.

For the first time, quantitative assessments of the degree of interaction during the formation of solid solutions according to a type VI reaction between Al 2 O 3 and nickel in a sintered material at one temperature (1673 K) were carried out by O. Kubashevsky. A detailed development of a method for thermodynamic assessment of the interaction of refractory oxides and a metal matrix of dispersion-strengthened materials was carried out by E.I. Mozzhukhin, whose calculation results were satisfactorily confirmed by chemical analysis of the A1 2 O 3 - Mo and A1 2 O 3 - Nb systems after their sintering at temperatures (0.6-0.8) of the matrix metal.

A type VI reaction can be accepted as the basis for thermodynamic calculations if the following conditions are met: the presence of at least a small solubility of oxygen and Me" in the welded metal Me"; no change in the stoichiometric composition of the oxide; no possibility of transition of the oxide participating in the reaction to lower oxides, lack of possibility of solubility of the welded metal in Me "" m O n.

Failure to fulfill the first condition deprives the equation in question of meaning: the second - leads to a reaction of type V; third - type VI reactions; fourth - necessitates the addition of reaction equation VI with another one, taking into account the formation of a solid solution Me" in and Me"" m O n of their joint solution.

In contrast to the reactions of types I, II, IV, V discussed above, for which the concept of thermodynamic equilibrium is not applicable and the direction of flow (from left to right or right to left) is entirely determined by the sign
, the type VI reaction proceeds from left to right and the completeness of its occurrence is determined by the equilibrium constant equal to the product of the activities of oxygen and Me "" in the welded metal Me". For dilute solutions, the activities can be taken equal to the concentration (mole fraction) and, using the law of mass action for a type VI reaction , determine their value, i.e. the equilibrium concentration of dissolved elements in the solid solution based on the metal being welded.The values ​​found will characterize the equilibrium degree of interaction of the materials being welded.

Thermodynamic calculation of a type VI reaction using the example of the ZnS-Me system, outlining the methodological features, is given in the work. The results of this calculation are, to a first approximation, also applicable to a similar ZnO-Me system, which is of particular interest when analyzing the weldability of zinc ferrites.

The calculation is based on the reaction of interaction with copper:

ZnS TV = Cu + [S] Cu (7.29)

The calculation results showed that when zinc sulfide interacts with copper, it is thermodynamically possible to dissolve up to 0.086 at. % sulfur, which is one and a half orders of magnitude higher than the solubility limit of sulfur in copper at this temperature (0.004 at.%), i.e. higher than can be contained in a saturated solid solution in equilibrium with lower copper sulfide. It follows that when ZnS interacts with copper, the formation of a certain amount of copper sulfide Cu 2 S is thermodynamically possible.

Consequently, thermodynamic calculation of interaction with copper according to the method of E.I. Mozzhukhin using equation (7.29) gives only a qualitative result. This technique is applicable for systems in which the difference in the Gibbs energies of formation of the refractory oxide and the matrix metal oxide is on the order of 400 kJ/g oxygen atom; in the sulfide systems under consideration, this value is much less.

To obtain quantitative results, further development of this technique is outlined below.

3

1 Moscow State Technical University named after. N.E. Bauman

2 First Moscow State medical University them. THEM. Sechenov

3 Moscow Pedagogical State University

The issues of etching oxide deposits from the surface of steels containing cobalt and iron have always been of practical importance and have been relevant. Having studied a large amount of material on this issue, the authors state that some aspects of the problem have not yet been fully studied (these include the influence of the characteristics of electrolyte solutions, identifying the mechanism of action of these factors). Cobalt and iron oxides are widely used as catalysts for various chemical processes (oxidation of methane and carbon monoxide, dehydrogenation of paraffins, etc.). Their properties depend on the characteristics of the surface, which determines the kinetics of oxide dissolution. Conducted experimental studies on the effect of mineral acids (in particular, H2SO4) on the rate of heterogeneous reaction (Co3O4 and Fe3O4 in an acidic medium) revealed the nature of the limiting stage, which consists in the formation of surface compounds of the type - and their subsequent transition to the electrolyte solution. A systematic analysis of oxide dissolution curves has also been developed to calculate kinetic parameters: activation energy and reaction orders for hydrogen ions and sulfate ions.

cobalt oxide

iron oxide

kinetics

dissolution

modeling

Barton–Stransky model

Hougen–Watson method

1. Bokshtein B.S., Mendelev M.I., Pokhvisnev Yu.V. Physical chemistry: thermodynamics and kinetics. – M.: Publishing house “MISIS”, 2012. – 258 p.

2. Butler J. Ionic equilibria. – L.: Chemistry, 1973. – 448 p.

3. Delmon B. Kinetics of heterogeneous reactions. – M.: Mir, 1972. – 555 p.

4. Barre P. Kinetics of heterogeneous processes. – M.: Mir, 1976. – 400 p.

5. Kiselev M.Yu. Mechanism and kinetics of pyrite dissolution by electrochemical chlorination // Izvestia Higher educational institutions. Mining magazine. – 2010. – No. 4. – P. 101–104.

6. Korzenshtein N.M., Samuylov E.V. Volumetric condensation in heterogeneous reactions // Colloid Journal. – 2013. – T. 75, No. 1. – 84 p.

7. Kolesnikov V.A., Kapustin V.A., Kapustin Yu.I., Isaev M.K., Kolesnikov A.V. Metal oxides – promising materials for electrochemical processes // Glass and Ceramics. – 2016. – No. 12. – P. 23–28.

8. Yakusheva E.A., Gorichev I.G., Atanasyan T.K., Izotov A.D. Study of the kinetics of dissolution of cobalt oxides (Co3O4, Co2O3) at various concentrations of H2SO4, HCl, EDTA and pH // Volgograd: Abstracts of XIX Mend. Congress on General and Applied Chemistry. – 2011. – T. 3 – P. 366.

9. Yakusheva E.A., Gorichev I.G., Atanasyan T.K., Layner Yu.A. Kinetics of dissolution of cobalt oxides in acidic media // Metals. – 2010. – No. 2. – P. 21–27.

10. Yakusheva E.A., Gorichev I.G., Atanasyan T.K., Plakhotnaya O.N., Goryacheva V.N. Modeling of kinetic processes of dissolution of cobalt and copper oxides in sulfuric acid // Bulletin of MSTU im. N.E. Bauman. Ser. Natural Sciences. – 2017. – No. 3. – pp. 124–134.

The conducted experimental studies of the dissolution of oxide phases make it possible to describe in detail the processes of the behavior of the solid phase in an acidic environment, to explain the phenomena occurring on the surface of oxides, taking into account their acid-base characteristics and the dissolution mechanism, and to simulate topochemical reactions.

Purpose of the study consists of studying and modeling the process of dissolution of Co3O4 and Fe3O4 in sulfuric acid.

Materials and research methods

For research, samples weighing 500 mg with d = 80÷100 µm were taken. Identification of oxides was carried out by X-ray diffraction, IR and thermal analyses.

To elucidate the mechanism of dissolution of solid samples of metal oxides in acidic media, the experiment was carried out in an instrument (a thermostated reactor with a volume of 0.5 l) to study the kinetics of dissolution of solid samples, excluding the influence of any uncontrolled factors on the phenomenon being studied. The experimental temperature was 363 K. The experiment was carried out at various pH values ​​and concentrations of mineral acid.

At certain time intervals, samples of the liquid phase were taken from the reaction vessel using a glass Schott filter. The concentration of cobalt ions was determined spectrophotometrically (UF-3100 spectrophotometer) using ammonium thiocyanate, and iron - using o-phenanthroline.

The obtained experimental data on the effect of acid concentration on the rate of dissolution of cobalt oxide Co3O4 and Fe3O4 are presented in Fig. 1 (dots - experimental data, lines - simulation results). The solute fraction a was calculated using the equation: a = Dt/D∞.

Rice. 1. a) dependence of the proportion of dissolved Co3O4 oxide on time at different concentrations of sulfuric acid (mol/l): 1 - 10.0; 2 - 5.93; 3 - 2.97; 4 - 1.0; 5 - 0.57; 6 - 0.12; T = 363.2 K; b) dependence of the proportion of dissolved Fe3O4 oxide on time at different concentrations of sulfuric acid (mol/l): 1 - 10.3; 2 - 7.82; 3 - 3.86; 4 - 2.44; T = 293 K

Research results and discussion

Calculation of kinetic parameters. An analysis of experimental kinetic data was carried out using the equations of heterogeneous kinetics, which made it possible to determine the orders of reactions for various ions (ni), the specific rate of dissolution (Wi), its dependence on the concentration of the solution, as well as the activation energies of reactions (Ea).

The kinetics of heterogeneous reactions is based on the mandatory consideration of changes in the surface of particles during the dissolution process over time; in addition, as a rule, heterogeneous reactions are characterized by a constant rate over time (1).

In this case, the rate of oxide dissolution can be represented by the equation:

where Wi is the specific dissolution rate; f(α) is a function that takes into account how the oxide surface changes over time.

To clarify the mechanism of dissolution and model this phenomenon, we used the Barton-Stransky model (2):

, (2)

where A is a constant. Its value is directly proportional to the number of active centers on the surface of one oxide particle.

To find the values ​​of the variables W and A, methods of nonlinear regression analysis and computer program MathCad.

Table 1

Specific rate of dissolution of oxides Co3O4 and Fe3O4 depending on the concentration of H2SO4

From the data in the table and Fig. 2 (dots - experimental data, lines - the result of modeling according to equation (3)) it follows that cobalt oxide Co3O4 dissolves faster in sulfuric acid than iron oxide Fe3O4. The reaction order in terms of hydrogen ions for the two oxides is approximately 0.5. (all results are based on the Barton-Stransky model).

Rice. 2. a) dependence of the logarithm of speed (log W) on the logarithm of concentration (log C(H2SO4)) when dissolving Co3O4 in sulfuric acid; b) dependence of the logarithm of speed (log W) on the logarithm of concentration (log C(H2SO4)) when Fe3O4 is dissolved in sulfuric acid

The data obtained make it possible to describe the relationship between the specific dissolution rate of Co3O4 and Fe3O4 oxides and the H2SO4 concentration by the generalized equation

, (3)

where ≡, W0 is the dissolution rate constant, K1, K2 are constants.

Modeling the mechanism of dissolution of cobalt and iron oxides in inorganic acid. The dissolution of oxides in acids occurs on surface defects of the crystal lattice, the so-called active centers of oxide dissolution that have adsorbed H+ ions and H+...A- ion pairs.

The Hougen-Watson method makes it possible to simulate the effect of pH and acid concentration on the rate of dissolution of oxides.

In this case, the dissolution rate of cobalt and iron oxides will be expressed by the equation:

Presumably, particles of metal hydroxo complexes of the same composition as those present in the solution are formed on the surface of the oxides. To calculate the concentration of hydroxo complexes, we used material balance equations in hydrolysis reactions for hydrogen, cobalt and iron ions; hydrolysis equations for all stages to calculate hydrolysis constants. The Hougen-Watson method assumes that the dependence of the ion concentration on the surface of the oxides and in the solution obeys the Langmuir isotherm, which makes it possible to relate the surface and volume concentrations of ions (equation (5)).

The dependence of the specific dissolution rate of cobalt oxides Co3O4 and Fe3O4 in dilute sulfuric acid is expressed by equations (5-7).

The concentration of ions and can be expressed in terms of the total concentration of Co3+ and Fe3+ ions, if their content in the solution is established. In this case and . Then the speed is

If we simulate the process of oxide dissolution and assume that the ions act as surface-active particles, then the dependence of the process speed on the ion concentration will look as follows (a1 is the number of ions in the solution).

§ 1 Oxide and its characteristics

While studying the chemical properties of oxygen, we became familiar with oxidation reactions and oxides. Oxides, for example, include substances with the following formulas: Na2O, CuO, Al2O3, SiO2, P2O5, SO3, Mn2O7.

So, all oxides are characterized in composition by three common features: any oxide is a complex substance, consisting of atoms of two chemical elements, one of the elements is oxygen.

All these features can be expressed by the general formula ExOy, in which E are atoms chemical element, which formed the oxide, O - oxygen atoms; x, y are indices indicating the number of atoms of the elements forming the oxide.

There are a lot of oxides. Almost all simple substances form oxides when oxidized. Atoms of many elements, exhibiting different valence values, participate in the formation of several oxides, for example, nitrogen corresponds to five oxides: nitrogen oxide (I) N2O, nitrogen oxide (II) NO, nitrogen oxide (III) N2O3, nitrogen oxide (IV) NO2, Nitric oxide (V) N2O5.

§ 2 Properties of oxides and their classification

Let's get acquainted with the properties of some oxides.

Carbon monoxide (IV) is a colorless, odorless gas with a slightly sour taste, turning into a solid white snow-like substance, bypassing the liquid state at - 780C, soluble in water.

Hydrogen oxide is water, under normal conditions it is a colorless liquid with a boiling point of 1000C.

Calcium oxide is a white solid whose melting point is 26270C; when mixed with water, it actively interacts with it.

Iron (III) oxide is a red-brown solid that melts at 15620C and is insoluble in water.

Let's pass carbon monoxide (IV) through water and add a few drops of litmus to the resulting solution. Litmus will change color from blue to red, therefore, when carbon monoxide (IV) reacts with water, an acid is formed. The reaction equation is as follows: CO2 + H2O → H2CO3. As a result of the reaction, carbonic acid was formed. Similarly, oxides of other nonmetals interact with water to form acids. Therefore, non-metal oxides are called acidic. Metal oxides with a valency greater than IV are also considered acidic, for example, vanadium (V) oxide V2O5, chromium (VI) oxide CrO3, manganese (VII) oxide Mn2O7.

Place some white calcium oxide powder in a test tube with water and add a few drops of phenolphthalein to the resulting slightly cloudy solution. Phenolphthalein changes color from colorless to crimson, which indicates the appearance of a base in the test tube. CaO + H2O → Ca(OH)2. As a result of the reaction, a base was formed - calcium hydroxide. Metal oxides whose valency is not more than III are called basic.

Metals exhibiting valence III and IV, and sometimes II, form amphoteric oxides. These oxides differ from others in their chemical properties. We will get to know them in more detail later, but for now we will focus our attention on acidic and basic oxides.

§ 3 Dissolution of oxides in water

Many acids and bases can be prepared by dissolving the corresponding oxides in water.

The dissolution of oxides in water is a chemical process accompanied by the formation of new chemical compounds - acids and bases.

For example, when sulfur oxide (VI) is dissolved in water, sulfuric acid is formed: SO3 + H2O → H2SO4. And when phosphorus oxide (V) is dissolved, phosphoric acid is formed: P2O5 + 3H2O → 2H3PO4. When sodium oxide is dissolved, a base is formed - sodium hydroxide: Na2O + H2O → 2NaOH, when barium oxide is dissolved - barium hydroxide: BaO + H2O → Ba(OH)2.

The names of oxide groups reflect their relationship with other classes of inorganic compounds: most acidic oxides correspond to acids, and almost all basic oxides correspond to bases.

However, not all oxides are soluble. Thus, most basic oxides are insoluble, and the only exceptions are oxides formed by elements of the main subgroups of the first and second groups of the periodic table of elements.

Most acid oxides, on the contrary, are soluble in water. The exception here is, for example, silicon (IV) oxide - SiO2. This substance is well known to everyone. Silicon oxide forms the basis of river sand and many minerals, including rare and very beautiful ones: rock crystal, amethyst, citrine, jasper. Many acidic oxides formed by metals are slightly soluble or insoluble.

If the oxides do not dissolve in water, then the corresponding acids and bases are obtained in other ways (indirectly), which we will get acquainted with later.

List of used literature:

1. NOT. Kuznetsova. Chemistry. 8th grade. Textbook for general education institutions. – M. Ventana-Graf, 2012.

increase

solubility of oxides and

hydroxides

Subgroup

When dissolving, ionic oxides enter into a chemical interaction with water, forming the corresponding hydroxides:

Na 2 O + H 2 O → 2NaOH

CaO + H 2 O → Ca(OH) 2

very strong

basic oxide base

Hydroxides of alkali and alkaline earth metals are strong bases and completely dissociate in water into metal cations and hydroxide ions:

NaOH Na + + OH –

Since the concentration of OH - ions increases, solutions of these substances have a highly alkaline environment (pH>>7); they are called alkalis.

Second group highly soluble in water oxides and their corresponding hydroxy compounds – molecular oxides and acids with covalent type of chemical bonds. These include compounds of typical nonmetals in the highest oxidation state and some d-metals in the oxidation state: +6, +7. Soluble molecular oxides (SO 3 , N 2 O 5 , Cl 2 O 7 , Mn 2 O 7 ) react with water to form the corresponding acids:

SO 3 + H 2 O H 2 SO 4

sulfur oxide (VI) sulfuric acid

strong acid strong acid

N2O5 + H2O2HNO3

nitric oxide (V) nitric acid

Mn 2 O 7 + H 2 O 2HMnO 4

manganese(VII) oxide manganese acid

Strong acids (H 2 SO 4, HNO 3, HClO 4, HClO 3, HMnO 4) in solutions completely dissociate into H + cations and acid residues:

Stage 2: H 2 PO 4 – H + + HPO 4 2–

K 2 =(=6.2∙10 –8;

Stage 3: HPO 4 2– H + + PO 4 3–

K 3 =()/=4.4∙10 –13 ,

where K1, K2, K3 are the dissociation constants of orthophosphoric acid, respectively, for the first, second and third stages.

The dissociation constant (Appendix Table 1) characterizes the strength of the acid, i.e. its ability to decompose (dissociate) into ions in a given solvent at a given temperature. The greater the dissociation constant, the more the equilibrium is shifted towards the formation of ions, the stronger the acid, i.e. In the first stage, the dissociation of phosphoric acid is better than in the second and, accordingly, in the third stage.

Moderately soluble oxides of sulfur (IV), carbon (IV), nitrogen (III), etc. form corresponding weak acids in water, which partially dissociate.

CO 2 + H 2 O H 2 CO 3 H + + HCO 3 –

SO 2 + H 2 O H 2 SO 3 H + + HSO 3 –

N 2 O 3 + H 2 O 2HNO 2 H + + NO 2 –

weak-weak

acidic acids

Neutralization reaction

The neutralization reaction can be expressed by the following scheme:

H 2 O

(base or (acid or acids-

basic oxide)

5.3.1. Properties of basic compounds exhibit oxides and hydroxides of s-metals (exception Be), d-metals in the oxidation state (+1, +2) (exception Zn), some p-metals (see Fig. 3).

										VIIIA
I A	II A				IIIA	IVA	V.A.	VIA	VIIA
Li	Be				B	C	N	O	F

Diagonal similarity Al Zn Ge Insoluble: usually basic Amphoteric oxides Weak acid Oxides dissolve to form acids

Rice. 3. Acid-base properties of oxides and their corresponding hydroxy compounds

A characteristic property of basic compounds is their ability to react with acids, acidic or amphoteric oxides to form salts, for example:

KOH + HCl KCl + H 2 O

Ba(OH) 2 + CO 2 BaCO 3 + H 2 O

2NaO + Al 2 O 3 2NaAlO 2 + H 2 O

Depending on the number of protons that can be added to the base, there are monoacid bases (for example, LiOH, KOH, NH 4 OH), diacid bases, etc.

For polyacid bases, the neutralization reaction can proceed in stages with the formation of first basic and then intermediate salts.

Me(OH) 2 MeOHCl MeCl 2

hydroxide NaOH basic NaOH medium

metal salt salt

For example:

Stage 1: Co(OH) 2 + HCl CoOHCl + H 2 O

hydroxocobalt(II)

(basic salt)

Stage 2: Co(OH)Cl + HCl CoCl 2 + H 2 O

cobalt(II)

(medium salt)

5.3.2. Properties of acid compounds exhibit oxides and acids of nonmetals, as well as d-metals in the oxidation state (+5, +6, +7) (see Fig. 3).

A characteristic property is their ability to interact with bases, basic and amphoteric oxides to form salts, for example:

2HNO 3 + Cu(OH) 2 → Cu(NO 3) 2 + 2H 2 O

2HCl + CaO → CaCl 2 + H 2 O

H 2 SO 4 + ZnO → ZnSO 4 + H 2 O

CrO 3 + 2NaOH → Na 2 CrO 4 + H 2 O

Based on the presence of oxygen in their composition, acids are divided into oxygen-containing(for example, H 2 SO 4, HNO 3) and oxygen-free(HBr, H 2 S). Based on the number of hydrogen atoms contained in an acid molecule that can be replaced by metal atoms, there are monobasic acids (for example, hydrogen chloride HCl, nitrous acid HNO 2), dibasic (sulphurous H 2 SO 3, coal H 2 CO 3), tribasic (orthophosphoric H 3 PO 4) etc.

Polybasic acids are neutralized stepwise with the formation of initially acidic and then medium salts:

H 2 X NaHX Na 2 X

polybasic acidic medium

acid salt salt

For example, orthophosphoric acid can form three types of salts depending on the quantitative ratio of the acid and alkali taken:

a) NaOH + H 3 PO 4 → NaH 2 PO 4 + H 2 O;

1:1 dihydrogen phosphate

b) 2NaOH + H 3 PO 4 → Na 2 HPO 4 + 2H 2 O;

2:1 hydrogen phosphate

c) 3NaOH + H 3 PO 4 → Na 3 PO 4 + 3H 2 O.

3:1 orthophosphate

5.3.3. Amphoteric oxides and hydroxides form Be, p-metals located near the “amphoteric diagonal” (Al, Ga, Sn, Pb), as well as d-metals in oxidation states (+3, +4) and Zn (+2) (see Fig. 3 ).

Slightly dissolving, amphoteric hydroxides dissociate both basic and acidic:

2H + + 2– Zn(OH) 2 Zn 2+ + 2OH –

Therefore, amphoteric oxides and hydroxides can react with both acids and bases. When interacting with stronger acids, amphoteric compounds exhibit the properties of bases.

ZnO + SO 3 → ZnSO 4 + H 2 O

acid

Zn(OH) 2 + H 2 SO 4 → ZnSO 4 + H 2 O

basic acid

connections

When interacting with strong bases, amphoteric compounds exhibit the properties of acids, forming the corresponding salts. The composition of the salt depends on the reaction conditions. When fused, simple “dehydrated” salts are formed.

2NaOH + Zn(OH) 2 → Na 2 ZnO 2 + H 2 O

acid base sodium zincate

compound

2NaOH + ZnO → Na 2 ZnO 2 + H 2 O

Complex salts are formed in aqueous solutions of alkalis:

2NaOH + Zn(OH) 2 → Na 2

(aqueous tetrahydroxozincate