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Acids can be classified based on different criteria:

1) The presence of oxygen atoms in the acid

2) Basicity of acid

The basicity of an acid is the number of “mobile” hydrogen atoms in its molecule, capable of being split off from the acid molecule during dissociation in the form of hydrogen cations H +, and also replaced by metal atoms:

4) Solubility

5) Stability

7) Oxidizing properties

Chemical properties of acids

1. Ability to dissociate

Acids dissociate in aqueous solutions into hydrogen cations and acid residues. As already mentioned, acids are divided into well-dissociating (strong) and low-dissociating (weak). When writing the dissociation equation for strong monobasic acids, either one right-pointing arrow () or an equal sign (=) is used, which shows the virtual irreversibility of such dissociation. For example, the dissociation equation for strong hydrochloric acid can be written in two ways:

or in this form: HCl = H + + Cl -

or in this way: HCl → H + + Cl -

In fact, the direction of the arrow tells us that the reverse process of combining hydrogen cations with acidic residues (association) practically does not occur in strong acids.

If we want to write the dissociation equation for a weak monoprotic acid, we must use two arrows in the equation instead of the sign. This sign reflects the reversibility of the dissociation of weak acids - in their case, the reverse process of combining hydrogen cations with acidic residues is strongly pronounced:

CH 3 COOH CH 3 COO — + H +

Polybasic acids dissociate stepwise, i.e. Hydrogen cations are separated from their molecules not simultaneously, but one by one. For this reason, the dissociation of such acids is expressed not by one, but by several equations, the number of which is equal to the basicity of the acid. For example, the dissociation of tribasic phosphoric acid occurs in three steps with the alternating separation of H + cations:

H 3 PO 4 H + + H 2 PO 4 —

H 2 PO 4 - H + + HPO 4 2-

HPO 4 2- H + + PO 4 3-

It should be noted that each subsequent stage of dissociation occurs to a lesser extent than the previous one. That is, H 3 PO 4 molecules dissociate better (to a greater extent) than H 2 PO 4 - ions, which, in turn, dissociate better than HPO 4 2- ions. This phenomenon is associated with an increase in the charge of acidic residues, as a result of which the strength of the bond between them and positive H + ions increases.

The exception to polybasic acids is sulfuric acid. Since this acid dissociates well in both stages, it is permissible to write the equation of its dissociation in one stage:

H 2 SO 4 2H + + SO 4 2-

2. Interaction of acids with metals

The seventh point in the classification of acids is their oxidizing properties. It was stated that acids are weak oxidizing agents and strong oxidizing agents. The vast majority of acids (almost all except H 2 SO 4 (conc.) and HNO 3) are weak oxidizing agents, since they can only exhibit their oxidizing ability due to hydrogen cations. Such acids can oxidize only those metals that are in the activity series to the left of hydrogen, and the salt of the corresponding metal and hydrogen are formed as products. For example:

H 2 SO 4 (diluted) + Zn ZnSO 4 + H 2

2HCl + Fe FeCl 2 + H 2

As for strong oxidizing acids, i.e. H 2 SO 4 (conc.) and HNO 3, then the list of metals on which they act is much wider, and it includes all metals before hydrogen in the activity series, and almost everything after. That is, concentrated sulfuric acid and Nitric acid of any concentration, for example, will oxidize even low-active metals such as copper, mercury, and silver. The interaction of nitric acid and concentrated sulfuric acid with metals, as well as some other substances, due to their specificity, will be discussed separately at the end of this chapter.

3. Interaction of acids with basic and amphoteric oxides

Acids react with basic and amphoteric oxides. Silicic acid, since it is insoluble, does not react with low-active basic oxides and amphoteric oxides:

H 2 SO 4 + ZnO ZnSO 4 + H 2 O

6HNO 3 + Fe 2 O 3 2Fe(NO 3) 3 + 3H 2 O

H 2 SiO 3 + FeO ≠

4. Interaction of acids with bases and amphoteric hydroxides

HCl + NaOH H 2 O + NaCl

3H 2 SO 4 + 2Al(OH) 3 Al 2 (SO 4) 3 + 6H 2 O

5. Interaction of acids with salts

This reaction occurs if a precipitate, gas, or a significantly weaker acid is formed than the one that reacts. For example:

H 2 SO 4 + Ba(NO 3) 2 BaSO 4 ↓ + 2HNO 3

CH 3 COOH + Na 2 SO 3 CH 3 COONa + SO 2 + H 2 O

HCOONa + HCl HCOOH + NaCl

6. Specific oxidative properties of nitric and concentrated sulfuric acids

As mentioned above, nitric acid in any concentration, as well as sulfuric acid exclusively in a concentrated state, are very strong oxidizing agents. In particular, unlike other acids, they oxidize not only metals that are located before hydrogen in the activity series, but also almost all metals after it (except platinum and gold).

For example, they are capable of oxidizing copper, silver and mercury. However, one should firmly grasp the fact that a number of metals (Fe, Cr, Al), despite the fact that they are quite active (available before hydrogen), nevertheless do not react with concentrated HNO 3 and concentrated H 2 SO 4 without heating due to the phenomenon of passivation - on the surface of such metals, protective film from solid oxidation products, which does not allow molecules of concentrated sulfuric and concentrated nitric acids to penetrate deep into the metal for the reaction to occur. However, with strong heating, the reaction still occurs.

In the case of interaction with metals, the obligatory products are always the salt of the corresponding metal and the acid used, as well as water. A third product is also always isolated, the formula of which depends on many factors, in particular, such as the activity of metals, as well as the concentration of acids and the reaction temperature.

The high oxidizing ability of concentrated sulfuric and concentrated nitric acids allows them to react not only with practically all metals of the activity series, but even with many solid non-metals, in particular with phosphorus, sulfur, and carbon. The table below clearly shows the products of the interaction of sulfuric and nitric acids with metals and non-metals depending on the concentration:

7. Reducing properties of oxygen-free acids

All oxygen-free acids (except HF) can exhibit reducing properties due to the chemical element included in the anion under the action of various oxidizing agents. For example, all hydrohalic acids (except HF) are oxidized by manganese dioxide, potassium permanganate, and potassium dichromate. In this case, halide ions are oxidized to free halogens:

4HCl + MnO 2 MnCl 2 + Cl 2 + 2H 2 O

18HBr + 2KMnO 4 2KBr + 2MnBr 2 + 8H 2 O + 5Br 2

14НI + K 2 Cr 2 O 7 3I 2 ↓ + 2Crl 3 + 2KI + 7H 2 O

Among all hydrohalic acids, hydroiodic acid has the greatest reducing activity. Unlike other hydrohalic acids, even ferric oxide and salts can oxidize it.

6HI ​​+ Fe 2 O 3 2FeI 2 + I 2 ↓ + 3H 2 O

2HI + 2FeCl 3 2FeCl 2 + I 2 ↓ + 2HCl

Hydrogen sulfide acid H 2 S also has high reducing activity. Even an oxidizing agent such as sulfur dioxide can oxidize it.

Acids are complex substances whose molecules include hydrogen atoms that can be replaced or exchanged for metal atoms and an acid residue.

Based on the presence or absence of oxygen in the molecule, acids are divided into oxygen-containing(H 2 SO 4 sulfuric acid, H 2 SO 3 sulfurous acid, HNO 3 nitric acid, H 3 PO 4 phosphoric acid, H 2 CO 3 carbonic acid, H 2 SiO 3 silicic acid) and oxygen-free(HF hydrofluoric acid, HCl hydrochloric acid ( hydrochloric acid), HBr hydrobromic acid, HI hydroiodic acid, H 2 S hydrosulfide acid).

Depending on the number of hydrogen atoms in the acid molecule, acids are monobasic (with 1 H atom), dibasic (with 2 H atoms) and tribasic (with 3 H atoms). For example, nitric acid HNO 3 is monobasic, since its molecule contains one hydrogen atom, sulfuric acid H 2 SO 4 – dibasic, etc.

There are very few inorganic compounds containing four hydrogen atoms that can be replaced by a metal.

The part of an acid molecule without hydrogen is called an acid residue.

Acidic residues may consist of one atom (-Cl, -Br, -I) - these are simple acidic residues, or they may consist of a group of atoms (-SO 3, -PO 4, -SiO 3) - these are complex residues.

In aqueous solutions, during exchange and substitution reactions, acidic residues are not destroyed:

H 2 SO 4 + CuCl 2 → CuSO 4 + 2 HCl

The word anhydride means anhydrous, that is, an acid without water. For example,

H 2 SO 4 – H 2 O → SO 3. Anoxic acids do not have anhydrides.

Acids get their name from the name of the acid-forming element (acid-forming agent) with the addition of the endings “naya” and less often “vaya”: H 2 SO 4 - sulfuric; H 2 SO 3 – coal; H 2 SiO 3 – silicon, etc.

The element can form several oxygen acids. In this case, the indicated endings in the names of acids will be when the element exhibits a higher valence (the acid molecule contains a high content of oxygen atoms). If the element exhibits a lower valence, the ending in the name of the acid will be “empty”: HNO 3 - nitric, HNO 2 - nitrogenous.

Acids can be obtained by dissolving anhydrides in water. If the anhydrides are insoluble in water, the acid can be obtained by the action of another stronger acid on the salt of the required acid. This method is typical for both oxygen and oxygen-free acids. Oxygen-free acids are also obtained by direct synthesis from hydrogen and a non-metal, followed by dissolving the resulting compound in water:

H 2 + Cl 2 → 2 HCl;

H 2 + S → H 2 S.

Solutions of the resulting gaseous substances HCl and H 2 S are acids.

Under normal conditions, acids exist in both liquid and solid states.

Chemical properties of acids

Acid solutions act on indicators. All acids (except silicic) are highly soluble in water. Special substances - indicators allow you to determine the presence of acid.

Indicators are substances of complex structure. They change their color depending on their interaction with different chemicals. In neutral solutions they have one color, in solutions of bases they have another color. When interacting with an acid, they change their color: the methyl orange indicator turns red, and the litmus indicator also turns red.

Interact with bases with the formation of water and salt, which contains an unchanged acid residue (neutralization reaction):

H 2 SO 4 + Ca(OH) 2 → CaSO 4 + 2 H 2 O.

Interact with base oxides with the formation of water and salt (neutralization reaction). The salt contains the acid residue of the acid that was used in the neutralization reaction:

H 3 PO 4 + Fe 2 O 3 → 2 FePO 4 + 3 H 2 O.

Interact with metals. For acids to interact with metals, certain conditions must be met:

1. the metal must be sufficiently active in relation to acids (in the series of activity of metals it must be located before hydrogen). The further to the left a metal is in the activity series, the more intensely it interacts with acids;

2. the acid must be strong enough (that is, capable of donating hydrogen ions H +).

When chemical reactions of acid with metals occur, salt is formed and hydrogen is released (except for the interaction of metals with nitric and concentrated sulfuric acids):

Zn + 2HCl → ZnCl 2 + H 2 ;

Cu + 4HNO 3 → CuNO 3 + 2 NO 2 + 2 H 2 O.

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Oxygen-free:	Basicity	Name of salt
HCl - hydrochloric (hydrochloric)	monobasic	chloride
HBr - hydrobromic	monobasic	bromide
HI - hydroiodide	monobasic	iodide
HF - hydrofluoric (fluoric)	monobasic	fluoride
H 2 S - hydrogen sulfide	dibasic	sulfide
Oxygen-containing:
HNO 3 – nitrogen	monobasic	nitrate
H 2 SO 3 - sulfurous	dibasic	sulfite
H 2 SO 4 – sulfuric	dibasic	sulfate
H 2 CO 3 - coal	dibasic	carbonate
H 2 SiO 3 - silicon	dibasic	silicate
H 3 PO 4 - orthophosphoric	tribasic	orthophosphate

Salts – complex substances that consist of metal atoms and acidic residues. This is the most numerous class of inorganic compounds.

Classification. By composition and properties: medium, acidic, basic, double, mixed, complex

Medium salts are products of complete replacement of the hydrogen atoms of a polybasic acid with metal atoms.

Upon dissociation, only metal cations (or NH 4 +) are produced. For example:

Na 2 SO 4 ® 2Na + +SO

CaCl 2 ® Ca 2+ + 2Cl -

Acid salts are products of incomplete replacement of hydrogen atoms of a polybasic acid with metal atoms.

Upon dissociation, they produce metal cations (NH 4 +), hydrogen ions and anions of the acid residue, for example:

NaHCO 3 ® Na + + HCO « H + +CO .

Basic salts are products of incomplete replacement of OH groups - the corresponding base with acidic residues.

Upon dissociation, they give metal cations, hydroxyl anions and an acid residue.

Zn(OH)Cl ® + + Cl - « Zn 2+ + OH - + Cl - .

Double salts contain two metal cations and upon dissociation give two cations and one anion.

KAl(SO 4) 2 ® K + + Al 3+ + 2SO

Complex salts contain complex cations or anions.

Br ® + + Br - « Ag + +2 NH 3 + Br -

Na ® Na + + - « Na + + Ag + + 2 CN -

Genetic relationship between different classes of compounds

EXPERIMENTAL PART

Equipment and utensils: rack with test tubes, washing machine, alcohol lamp.

Reagents and materials: red phosphorus, zinc oxide, Zn granules, slaked lime powder Ca(OH) 2, 1 mol/dm 3 solutions of NaOH, ZnSO 4, CuSO 4, AlCl 3, FeCl 3, HСl, H 2 SO 4, universal indicator paper, solution phenolphthalein, methyl orange, distilled water.

Work order

1. Pour zinc oxide into two test tubes; add an acid solution (HCl or H 2 SO 4) to one and an alkali solution (NaOH or KOH) to the other and heat slightly on an alcohol lamp.

Observations: Does zinc oxide dissolve in an acid and alkali solution?

Write equations

Conclusions: 1.What type of oxide does ZnO belong to?

2. What properties do amphoteric oxides have?

Preparation and properties of hydroxides

2.1. Dip the tip of the universal indicator strip into the alkali solution (NaOH or KOH). Compare the resulting color of the indicator strip with the standard color scale.

Observations: Record the pH value of the solution.

2.2. Take four test tubes, pour 1 ml of ZnSO 4 solution into the first, CuSO 4 into the second, AlCl 3 into the third, and FeCl 3 into the fourth. Add 1 ml of NaOH solution to each test tube. Write observations and equations for the reactions occurring.

Observations: Does precipitation occur when alkali is added to a salt solution? Indicate the color of the sediment.

Write equations occurring reactions (in molecular and ionic form).

Conclusions: How can metal hydroxides be prepared?

2.3. Transfer half of the sediments obtained in experiment 2.2 to other test tubes. Treat one part of the sediment with a solution of H 2 SO 4 and the other with a solution of NaOH.

Observations: Does precipitate dissolution occur when alkali and acid are added to precipitates?

Write equations occurring reactions (in molecular and ionic form).

Conclusions: 1. What type of hydroxides are Zn(OH) 2, Al(OH) 3, Cu(OH) 2, Fe(OH) 3?

2. What properties do amphoteric hydroxides have?

Obtaining salts.

3.1. Pour 2 ml of CuSO 4 solution into a test tube and dip a cleaned nail into this solution. (The reaction is slow, changes on the surface of the nail appear after 5-10 minutes).

Observations: Are there any changes to the surface of the nail? What is being deposited?

Write the equation for the redox reaction.

Conclusions: Taking into account the range of metal stresses, indicate the method of obtaining salts.

3.2. Place one zinc granule in a test tube and add HCl solution.

Observations: Is there any gas evolution?

Write the equation

Conclusions: Explain this method of obtaining salts?

3.3. Pour some slaked lime powder Ca(OH) 2 into a test tube and add HCl solution.

Observations: Is there gas evolution?

Write the equation the reaction taking place (in molecular and ionic form).

Conclusion: 1. What type of reaction is the interaction between a hydroxide and an acid?

2.What substances are the products of this reaction?

3.5. Pour 1 ml of salt solutions into two test tubes: into the first - copper sulfate, into the second - cobalt chloride. Add to both test tubes drop by drop sodium hydroxide solution until precipitation forms. Then add excess alkali to both test tubes.

Observations: Indicate the changes in the color of precipitation in the reactions.

Write the equation the reaction taking place (in molecular and ionic form).

Conclusion: 1. As a result of what reactions are basic salts formed?

2. How can you convert basic salts to medium salts?

Test tasks:

1. From the listed substances, write down the formulas of salts, bases, acids: Ca(OH) 2, Ca(NO 3) 2, FeCl 3, HCl, H 2 O, ZnS, H 2 SO 4, CuSO 4, KOH
Zn(OH) 2, NH 3, Na 2 CO 3, K 3 PO 4.

2. Indicate the formulas of the oxides corresponding to the listed substances H 2 SO 4, H 3 AsO 3, Bi(OH) 3, H 2 MnO 4, Sn(OH) 2, KOH, H 3 PO 4, H 2 SiO 3, Ge( OH) 4 .

3. Which hydroxides are amphoteric? Write down reaction equations characterizing the amphotericity of aluminum hydroxide and zinc hydroxide.

4. Which of the following compounds will interact in pairs: P 2 O 5 , NaOH, ZnO, AgNO 3 , Na 2 CO 3 , Cr(OH) 3 , H 2 SO 4 . Write down equations for possible reactions.

Laboratory work No. 2 (4 hours)

Subject: Qualitative analysis of cations and anions

Target: master the technique of conducting qualitative and group reactions on cations and anions.

THEORETICAL PART

The main task of qualitative analysis is to establish chemical composition substances found in various objects (biological materials, medicines, food products, objects environment). IN this work The qualitative analysis of inorganic substances that are electrolytes is considered, i.e., essentially a qualitative analysis of ions. From the entire set of occurring ions, the most important in medical and biological terms were selected: (Fe 3+, Fe 2+, Zn 2+, Ca 2+, Na +, K +, Mg 2+, Cl -, PO, CO, etc. ). Many of these ions are part of various medicines and food products.

Not all are used in qualitative analysis possible reactions, but only those that are accompanied by a clear analytical effect. The most common analytical effects: the appearance of a new color, the release of gas, the formation of a precipitate.

There are two fundamentally different approaches to qualitative analysis: fractional and systematic . In systematic analysis, group reagents are necessarily used to separate the ions present into separate groups, and in some cases into subgroups. To do this, some of the ions are converted into insoluble compounds, and some of the ions are left in solution. After separating the precipitate from the solution, they are analyzed separately.

For example, the solution contains A1 3+, Fe 3+ and Ni 2+ ions. If this solution is exposed to excess alkali, a precipitate of Fe(OH) 3 and Ni(OH) 2 precipitates, and [A1(OH) 4 ] - ions remain in the solution. The precipitate containing iron and nickel hydroxides will partially dissolve when treated with ammonia due to the transition to 2+ solution. Thus, using two reagents - alkali and ammonia, two solutions were obtained: one contained [A1(OH) 4 ] - ions, the other contained 2+ ions and a Fe(OH) 3 precipitate. Using characteristic reactions, the presence of certain ions is then proven in solutions and in the precipitate, which must first be dissolved.

Systematic analysis is used mainly for the detection of ions in complex multicomponent mixtures. It is very labor-intensive, but its advantage lies in the easy formalization of all actions that fit into a clear scheme (methodology).

To carry out fractional analysis, only characteristic reactions are used. Obviously, the presence of other ions can significantly distort the results of the reaction (overlapping colors, unwanted precipitation, etc.). To avoid this, fractional analysis mainly uses highly specific reactions that give an analytical effect with a small number of ions. For successful reactions, it is very important to maintain certain conditions, in particular pH. Very often in fractional analysis it is necessary to resort to masking, that is, to convert ions into compounds that are not capable of producing an analytical effect with the selected reagent. For example, dimethylglyoxime is used to detect nickel ion. The Fe 2+ ion gives a similar analytical effect to this reagent. To detect Ni 2+, the Fe 2+ ion is transferred to a stable fluoride complex 4- or oxidized to Fe 3+, for example, with hydrogen peroxide.

Fractional analysis is used to detect ions in simpler mixtures. The analysis time is significantly reduced, but at the same time the experimenter is required to have a deeper knowledge of the patterns of chemical reactions, since it is quite difficult to take into account in one specific technique all possible cases of mutual influence of ions on the nature of the observed analytical effects.

In analytical practice, the so-called fractional-systematic method. With this approach, a minimum number of group reagents is used, which makes it possible to outline analysis tactics in general outline, which is then carried out using the fractional method.

According to the technique of conducting analytical reactions, reactions are distinguished: sedimentary; microcrystalscopic; accompanied by the release of gaseous products; conducted on paper; extraction; colored in solutions; flame coloring.

When carrying out sedimentary reactions, the color and nature of the precipitate (crystalline, amorphous) must be noted; if necessary, additional tests are carried out: the precipitate is checked for solubility in strong and weak acids, alkalis and ammonia, and an excess of the reagent. When carrying out reactions accompanied by the release of gas, its color and smell are noted. In some cases, additional tests are carried out.

For example, if the gas released is suspected to be carbon monoxide (IV), it is passed through an excess of lime water.

In fractional and systematic analyses, reactions during which a new color appears are widely used, most often these are complexation reactions or redox reactions.

In some cases, it is convenient to carry out such reactions on paper (droplet reactions). Reagents that do not decompose under normal conditions are applied to paper in advance. Thus, to detect hydrogen sulfide or sulfide ions, paper impregnated with lead nitrate is used [blackening occurs due to the formation of lead(II) sulfide]. Many oxidizing agents are detected using iodine starch paper, i.e. paper soaked in solutions of potassium iodide and starch. In most cases, the necessary reagents are applied to paper during the reaction, for example, alizarin for the A1 3+ ion, cupron for the Cu 2+ ion, etc. To enhance the color, extraction into an organic solvent is sometimes used. For preliminary tests, flame color reactions are used.

Classification of inorganic substances with examples of compounds

Now let's analyze the classification scheme presented above in more detail.

As we see, first of all, all inorganic substances are divided into simple And complex:

Simple substances These are substances that are formed by atoms of only one chemical element. For example, simple substances are hydrogen H2, oxygen O2, iron Fe, carbon C, etc.

Among simple substances there are metals, nonmetals And noble gases:

Metals formed by chemical elements located below the boron-astatine diagonal, as well as all elements located in side groups.

Noble gases formed by chemical elements of group VIIIA.

Nonmetals are formed respectively by chemical elements located above the boron-astatine diagonal, with the exception of all elements of side subgroups and noble gases located in group VIIIA:

The names of simple substances most often coincide with the names of the chemical elements whose atoms they are formed from. However, for many chemical elements the phenomenon of allotropy is widespread. Allotropy is the phenomenon when one chemical element capable of forming several simple substances. For example, in the case of the chemical element oxygen, the existence of molecular compounds with the formulas O 2 and O 3 is possible. The first substance is usually called oxygen in the same way as the chemical element whose atoms it is formed, and the second substance (O 3) is usually called ozone. Under simple substance carbon can mean any of its allotropic modifications, for example, diamond, graphite or fullerenes. The simple substance phosphorus can be understood as its allotropic modifications, such as white phosphorus, red phosphorus, black phosphorus.

Complex substances

Complex substances are substances formed by atoms of two or more chemical elements.

For example, complex substances are ammonia NH 3, sulfuric acid H 2 SO 4, slaked lime Ca(OH) 2 and countless others.

Among complex inorganic substances, there are 5 main classes, namely oxides, bases, amphoteric hydroxides, acids and salts:

Oxides - complex substances formed by two chemical elements, one of which is oxygen in the oxidation state -2.

The general formula of oxides can be written as E x O y, where E is the symbol of a chemical element.

Nomenclature of oxides

The name of the oxide of a chemical element is based on the principle:

For example:

Fe 2 O 3 - iron (III) oxide; CuO—copper(II) oxide; N 2 O 5 - nitric oxide (V)

You can often find information that the valency of an element is indicated in parentheses, but this is not the case. So, for example, the oxidation state of nitrogen N 2 O 5 is +5, and the valence, oddly enough, is four.

If a chemical element has a single positive oxidation state in compounds, then the oxidation state is not indicated. For example:

Na 2 O - sodium oxide; H 2 O - hydrogen oxide; ZnO - zinc oxide.

Oxides classification

Oxides, according to their ability to form salts when interacting with acids or bases, are divided accordingly into salt-forming And non-salt-forming.

There are few non-salt-forming oxides; they are all formed by nonmetals in the oxidation state +1 and +2. The list of non-salt-forming oxides should be remembered: CO, SiO, N 2 O, NO.

Salt-forming oxides, in turn, are divided into basic, acidic And amphoteric.

Basic oxides These are oxides that, when reacting with acids (or acid oxides), form salts. Basic oxides include metal oxides in the oxidation state +1 and +2, with the exception of the oxides BeO, ZnO, SnO, PbO.

Acidic oxides These are oxides that, when reacting with bases (or basic oxides), form salts. Acidic oxides are almost all oxides of non-metals with the exception of non-salt-forming CO, NO, N 2 O, SiO, as well as all metal oxides in high oxidation states (+5, +6 and +7).

Amphoteric oxides are called oxides that can react with both acids and bases, and as a result of these reactions they form salts. Such oxides exhibit a dual acid-base nature, that is, they can exhibit the properties of both acidic and basic oxides. Amphoteric oxides include metal oxides in the oxidation states +3, +4, as well as the oxides BeO, ZnO, SnO, and PbO as exceptions.

Some metals can form all three types of salt-forming oxides. For example, chromium forms the basic oxide CrO, the amphoteric oxide Cr 2 O 3 and the acidic oxide CrO 3.

As you can see, the acid-base properties of metal oxides directly depend on the degree of oxidation of the metal in the oxide: the higher the degree of oxidation, the more pronounced the acidic properties.

Grounds

Grounds - compounds with the formula Me(OH) x, where x most often equal to 1 or 2.

Classification of bases

Bases are classified according to the number of hydroxyl groups in one structural unit.

Bases with one hydroxo group, i.e. type MeOH is called monoacid bases, with two hydroxo groups, i.e. type Me(OH) 2, respectively, diacid etc.

Bases are also divided into soluble (alkalis) and insoluble.

Alkalies include exclusively hydroxides of alkali and alkaline earth metals, as well as thallium hydroxide TlOH.

Nomenclature of bases

The name of the foundation is based on the following principle:

For example:

Fe(OH) 2 - iron (II) hydroxide,

Cu(OH) 2 - copper (II) hydroxide.

In cases where the metal in complex substances has a constant oxidation state, it is not required to indicate it. For example:

NaOH - sodium hydroxide,

Ca(OH) 2 - calcium hydroxide, etc.

Acids

Acids - complex substances whose molecules contain hydrogen atoms that can be replaced by a metal.

The general formula of acids can be written as H x A, where H are hydrogen atoms that can be replaced by a metal, and A is the acidic residue.

For example, acids include compounds such as H2SO4, HCl, HNO3, HNO2, etc.

Classification of acids

According to the number of hydrogen atoms that can be replaced by a metal, acids are divided into:

- O base acids: HF, HCl, HBr, HI, HNO 3 ;

- d basic acids: H 2 SO 4, H 2 SO 3, H 2 CO 3;

- T rehobasic acids: H 3 PO 4 , H 3 BO 3 .

It should be noted that the number of hydrogen atoms in the case of organic acids most often does not reflect their basicity. For example, acetic acid with the formula CH 3 COOH, despite the presence of 4 hydrogen atoms in the molecule, is not tetra- but monobasic. The basicity of organic acids is determined by the number of carboxyl groups (-COOH) in the molecule.

Also, based on the presence of oxygen in the molecules, acids are divided into oxygen-free (HF, HCl, HBr, etc.) and oxygen-containing (H 2 SO 4, HNO 3, H 3 PO 4, etc.). Oxygen-containing acids are also called oxoacids.

You can read more about the classification of acids.

Nomenclature of acids and acid residues

The following list of names and formulas of acids and acid residues is a must-learn.

In some cases, a number of the following rules can make memorization easier.

As can be seen from the table above, the construction of systematic names of oxygen-free acids is as follows:

For example:

HF—hydrofluoric acid;

HCl—hydrochloric acid;

H 2 S is hydrosulfide acid.

The names of acidic residues of oxygen-free acids are based on the principle:

For example, Cl - - chloride, Br - - bromide.

The names of oxygen-containing acids are obtained by adding various suffixes and endings to the name of the acid-forming element. For example, if the acid-forming element in an oxygen-containing acid has the highest oxidation state, then the name of such an acid is constructed as follows:

For example, sulfuric acid H 2 S +6 O 4, chromic acid H 2 Cr +6 O 4.

All oxygen-containing acids can also be classified as acid hydroxides because they contain hydroxyl groups (OH). For example, this can be seen from the following graphical formulas of some oxygen-containing acids:

Thus, sulfuric acid can otherwise be called sulfur (VI) hydroxide, nitric acid - nitrogen (V) hydroxide, phosphoric acid - phosphorus (V) hydroxide, etc. In this case, the number in brackets characterizes the degree of oxidation of the acid-forming element. This version of the names of oxygen-containing acids may seem extremely unusual to many, but occasionally such names can be found in real KIMs of the Unified State Examination in Chemistry in tasks on the classification of inorganic substances.

Amphoteric hydroxides

Amphoteric hydroxides - metal hydroxides exhibiting a dual nature, i.e. capable of exhibiting both the properties of acids and the properties of bases.

Metal hydroxides in oxidation states +3 and +4 are amphoteric (as are oxides).

Also, as exceptions, amphoteric hydroxides include the compounds Be(OH) 2, Zn(OH) 2, Sn(OH) 2 and Pb(OH) 2, despite the oxidation state of the metal in them +2.

For amphoteric hydroxides of tri- and tetravalent metals, the existence of ortho- and meta-forms is possible, differing from each other by one water molecule. For example, aluminum(III) hydroxide can exist in the ortho form Al(OH)3 or the meta form AlO(OH) (metahydroxide).

Since, as already mentioned, amphoteric hydroxides exhibit both the properties of acids and the properties of bases, their formula and name can also be written differently: either as a base or as an acid. For example:

Salts

For example, salts include compounds such as KCl, Ca(NO 3) 2, NaHCO 3, etc.

The definition presented above describes the composition of most salts, however, there are salts that do not fall under it. For example, instead of metal cations, the salt may contain ammonium cations or its organic derivatives. Those. salts include compounds such as, for example, (NH 4) 2 SO 4 (ammonium sulfate), + Cl - (methyl ammonium chloride), etc.

Classification of salts

On the other hand, salts can be considered as products of the replacement of hydrogen cations H + in an acid with other cations, or as products of the replacement of hydroxide ions in bases (or amphoteric hydroxides) with other anions.

With complete replacement, the so-called average or normal salt. For example, with complete replacement of hydrogen cations in sulfuric acid with sodium cations, an average (normal) salt Na 2 SO 4 is formed, and with complete replacement of hydroxide ions in the base Ca (OH) 2 with acidic residues of nitrate ions, an average (normal) salt is formed Ca(NO3)2.

Salts obtained by incomplete replacement of hydrogen cations in a dibasic (or more) acid with metal cations are called acidic. Thus, when hydrogen cations in sulfuric acid are incompletely replaced by sodium cations, the acid salt NaHSO 4 is formed.

Salts that are formed by incomplete replacement of hydroxide ions in two-acid (or more) bases are called bases. O strong salts. For example, with incomplete replacement of hydroxide ions in the base Ca(OH) 2 with nitrate ions, a base is formed O clear salt Ca(OH)NO3.

Salts consisting of cations of two different metals and anions of acidic residues of only one acid are called double salts. So, for example, double salts are KNaCO 3, KMgCl 3, etc.

If a salt is formed by one type of cations and two types of acid residues, such salts are called mixed. For example, mixed salts are the compounds Ca(OCl)Cl, CuBrCl, etc.

There are salts that do not fall under the definition of salts as products of the replacement of hydrogen cations in acids with metal cations or products of the replacement of hydroxide ions in bases with anions of acidic residues. These are complex salts. For example, complex salts are sodium tetrahydroxozincate and tetrahydroxoaluminate with the formulas Na 2 and Na, respectively. Complex salts can most often be recognized among others by the presence of square brackets in the formula. However, you need to understand that in order for a substance to be classified as a salt, it must contain some cations other than (or instead of) H +, and the anions must contain some anions other than (or instead of) OH -. For example, the compound H2 does not belong to the class of complex salts, since when it dissociates from cations, only hydrogen cations H+ are present in the solution. Based on the type of dissociation, this substance should rather be classified as an oxygen-free complex acid. Likewise, the OH compound does not belong to salts, because this compound consists of cations + and hydroxide ions OH -, i.e. it should be considered a comprehensive foundation.

Nomenclature of salts

Nomenclature of medium and acid salts

The name of the middle and acid salts is built on the principle:

If the oxidation state of a metal in complex substances is constant, then it is not indicated.

The names of acid residues were given above when considering the nomenclature of acids.

For example,

Na 2 SO 4 - sodium sulfate;

NaHSO 4 - sodium hydrogen sulfate;

CaCO 3 - calcium carbonate;

Ca(HCO 3) 2 - calcium bicarbonate, etc.

Nomenclature of basic salts

The names of the main salts are based on the principle:

For example:

(CuOH) 2 CO 3 - copper (II) hydroxycarbonate;

Fe(OH) 2 NO 3 - iron (III) dihydroxonitrate.

Nomenclature of complex salts

The nomenclature of complex compounds is much more complicated, and to pass the Unified State Exam you do not need to know much about the nomenclature of complex salts.

You should be able to name complex salts obtained by reacting alkali solutions with amphoteric hydroxides. For example:

*The same colors in the formula and name indicate the corresponding elements of the formula and name.

Trivial names of inorganic substances

By trivial names we mean the names of substances that are not related, or weakly related, to their composition and structure. Trivial names are determined, as a rule, either historical reasons either physical or chemical properties connection data.

List of trivial names of inorganic substances that you need to know:

Na 3	cryolite
SiO2	quartz, silica
FeS 2	pyrite, iron pyrite
CaSO 4 ∙2H 2 O	gypsum
CaC2	calcium carbide
Al 4 C 3	aluminum carbide
KOH	caustic potassium
NaOH	caustic soda, caustic soda
H2O2	hydrogen peroxide
CuSO 4 ∙5H 2 O	copper sulfate
NH4Cl	ammonia
CaCO3	chalk, marble, limestone
N2O	laughing gas
NO 2	brown gas
NaHCO3	baking (drinking) soda
Fe3O4	iron scale
NH 3 ∙H 2 O (NH 4 OH)	ammonia
CO	carbon monoxide
CO2	carbon dioxide
SiC	carborundum (silicon carbide)
PH 3	phosphine
NH 3	ammonia
KClO3	Bertholet's salt (potassium chlorate)
(CuOH)2CO3	malachite
CaO	quicklime
Ca(OH)2	slaked lime
transparent aqueous solution of Ca(OH) 2	lime water
suspension of solid Ca(OH) 2 in its aqueous solution	lime milk
K2CO3	potash
Na 2 CO 3	soda ash
Na 2 CO 3 ∙10H 2 O	crystal soda
MgO	magnesia