Acids and their formulas. Density of various acids. Trivial names of inorganic substances

7. Acids. Salt. Relationship between classes of inorganic substances

7.1. Acids

Acids are electrolytes, upon the dissociation of which only hydrogen cations H + are formed as positively charged ions (more precisely, hydronium ions H 3 O +).

Another definition: acids are complex substances consisting of a hydrogen atom and acid residues (Table 7.1).

Table 7.1

Formulas and names of some acids, acid residues and salts

Under normal conditions, acids can be solids (H 3 PO 4, H 3 BO 3, H 2 SiO 3) and liquids (HNO 3, H 2 SO 4, CH 3 COOH). These acids can exist both individually (100% form) and in the form of diluted and concentrated solutions. For example, H 2 SO 4 , HNO 3 , H 3 PO 4 , CH 3 COOH are known both individually and in solutions.

A number of acids are known only in solutions. These are all hydrogen halides (HCl, HBr, HI), hydrogen sulfide H 2 S, hydrogen cyanide (hydrocyanic HCN), carbonic H 2 CO 3, sulfurous H 2 SO 3 acid, which are solutions of gases in water. For example, hydrochloric acid is a mixture of HCl and H 2 O, carbonic acid is a mixture of CO 2 and H 2 O. It is clear that using the expression “hydrochloric acid solution” is incorrect.

Most acids are soluble in water; silicic acid H 2 SiO 3 is insoluble. The overwhelming majority of acids have a molecular structure. Examples of structural formulas of acids:

In most oxygen-containing acid molecules, all hydrogen atoms are bonded to oxygen. But there are exceptions:

Acids are classified according to a number of characteristics (Table 7.2).

Table 7.2

Classification of acids

All general chemical properties of acids are due to the presence in their aqueous solutions of excess hydrogen cations H + (H 3 O +).

1. Due to the excess of H + ions, aqueous solutions of acids change the color of litmus violet and methyl orange to red (phenolphthalein does not change color and remains colorless). In an aqueous solution of weak carbonic acid, litmus is not red, but pink; a solution over a precipitate of very weak silicic acid does not change the color of the indicators at all.

2. Acids interact with basic oxides, bases and amphoteric hydroxides, ammonia hydrate (see Chapter 6).

Example 7.1. To carry out the transformation BaO → BaSO 4 you can use: a) SO 2; b) H 2 SO 4; c) Na 2 SO 4; d) SO 3.

Solution. The transformation can be carried out using H 2 SO 4:

BaO + H 2 SO 4 = BaSO 4 ↓ + H 2 O

BaO + SO 3 = BaSO 4

Na 2 SO 4 does not react with BaO, and in the reaction of BaO with SO 2 barium sulfite is formed:

BaO + SO 2 = BaSO 3

Answer: 3).

3. Acids react with ammonia and its aqueous solutions to form ammonium salts:

HCl + NH 3 = NH 4 Cl - ammonium chloride;

H 2 SO 4 + 2NH 3 = (NH 4) 2 SO 4 - ammonium sulfate.

4. Non-oxidizing acids react with metals located in the activity series up to hydrogen to form a salt and release hydrogen:

H 2 SO 4 (diluted) + Fe = FeSO 4 + H 2

2HCl + Zn = ZnCl 2 = H 2

The interaction of oxidizing acids (HNO 3, H 2 SO 4 (conc)) with metals is very specific and is considered when studying the chemistry of elements and their compounds.

5. Acids interact with salts. The reaction has a number of features:

a) in most cases, when a stronger acid reacts with a salt of a weaker acid, a salt of a weak acid and a weak acid are formed, or, as they say, a stronger acid displaces a weaker one. The series of decreasing strength of acids looks like this:

Examples of reactions occurring:

2HCl + Na 2 CO 3 = 2NaCl + H 2 O + CO 2

H 2 CO 3 + Na 2 SiO 3 = Na 2 CO 3 + H 2 SiO 3 ↓

2CH 3 COOH + K 2 CO 3 = 2CH 3 COOK + H 2 O + CO 2

3H 2 SO 4 + 2K 3 PO 4 = 3K 2 SO 4 + 2H 3 PO 4

Do not interact with each other, for example, KCl and H 2 SO 4 (diluted), NaNO 3 and H 2 SO 4 (diluted), K 2 SO 4 and HCl (HNO 3, HBr, HI), K 3 PO 4 and H 2 CO 3, CH 3 COOK and H 2 CO 3;

b) in some cases, a weaker acid displaces a stronger one from a salt:

CuSO 4 + H 2 S = CuS↓ + H 2 SO 4

3AgNO 3 (dil) + H 3 PO 4 = Ag 3 PO 4 ↓ + 3HNO 3.

Such reactions are possible when the precipitates of the resulting salts do not dissolve in the resulting dilute strong acids (H 2 SO 4 and HNO 3);

c) in the case of the formation of precipitates that are insoluble in strong acids, a reaction may occur between a strong acid and a salt formed by another strong acid:

BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl

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

AgNO 3 + HCl = AgCl↓ + HNO 3

Example 7.2. Indicate the row containing the formulas of substances that react with H 2 SO 4 (diluted).

1) Zn, Al 2 O 3, KCl (p-p); 3) NaNO 3 (p-p), Na 2 S, NaF; 2) Cu(OH) 2, K 2 CO 3, Ag; 4) Na 2 SO 3, Mg, Zn(OH) 2.

Solution. All substances of row 4 interact with H 2 SO 4 (dil):

Na 2 SO 3 + H 2 SO 4 = Na 2 SO 4 + H 2 O + SO 2

Mg + H 2 SO 4 = MgSO 4 + H 2

Zn(OH) 2 + H 2 SO 4 = ZnSO 4 + 2H 2 O

In row 1) the reaction with KCl (p-p) is not feasible, in row 2) - with Ag, in row 3) - with NaNO 3 (p-p).

Answer: 4).

6. Concentrated sulfuric acid behaves very specifically in reactions with salts. This is a non-volatile and thermally stable acid, therefore it displaces all strong acids from solid (!) salts, since they are more volatile than H2SO4 (conc):

KCl (tv) + H 2 SO 4 (conc.) KHSO 4 + HCl

2KCl (s) + H 2 SO 4 (conc) K 2 SO 4 + 2HCl

Salts formed by strong acids (HBr, HI, HCl, HNO 3, HClO 4) react only with concentrated sulfuric acid and only when in a solid state

Example 7.3. Concentrated sulfuric acid, unlike dilute one, reacts:

3) KNO 3 (tv);

Solution. Both acids react with KF, Na 2 CO 3 and Na 3 PO 4, and only H 2 SO 4 (conc.) react with KNO 3 (solid).

Answer: 3).

Methods for producing acids are very diverse.

Anoxic acids receive:

• by dissolving the corresponding gases in water:

HCl (g) + H 2 O (l) → HCl (p-p)

H 2 S (g) + H 2 O (l) → H 2 S (solution)

• from salts by displacement with stronger or less volatile acids:

FeS + 2HCl = FeCl 2 + H 2 S

KCl (tv) + H 2 SO 4 (conc) = KHSO 4 + HCl

Na 2 SO 3 + H 2 SO 4 Na 2 SO 4 + H 2 SO 3

Oxygen-containing acids receive:

• by dissolving the corresponding acidic oxides in water, while the degree of oxidation of the acid-forming element in the oxide and acid remains the same (with the exception of NO 2):

N2O5 + H2O = 2HNO3

SO 3 + H 2 O = H 2 SO 4

P 2 O 5 + 3H 2 O 2H 3 PO 4

• oxidation of non-metals with oxidizing acids:

S + 6HNO 3 (conc) = H 2 SO 4 + 6NO 2 + 2H 2 O

• by displacing a strong acid from a salt of another strong acid (if a precipitate insoluble in the resulting acids precipitates):

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

AgNO 3 + HCl = AgCl↓ + HNO 3

• by displacing a volatile acid from its salts with a less volatile acid.

For this purpose, non-volatile, thermally stable concentrated sulfuric acid is most often used:

NaNO 3 (tv) + H 2 SO 4 (conc.) NaHSO 4 + HNO 3

KClO 4 (tv) + H 2 SO 4 (conc.) KHSO 4 + HClO 4

• displacement of a weaker acid from its salts by a stronger acid:

Ca 3 (PO 4) 2 + 3H 2 SO 4 = 3CaSO 4 ↓ + 2H 3 PO 4

NaNO 2 + HCl = NaCl + HNO 2

K 2 SiO 3 + 2HBr = 2KBr + H 2 SiO 3 ↓

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 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 with respect 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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Acids- electrolytes, upon dissociation of which only H + ions are formed from positive ions:

HNO 3 ↔ H + + NO 3 - ;

CH 3 COOH↔ H + +CH 3 COO — .

All acids are classified into inorganic and organic (carboxylic), which also have their own (internal) classifications.

Under normal conditions, a significant amount of inorganic acids exist in a liquid state, some in a solid state (H 3 PO 4, H 3 BO 3).

Organic acids with up to 3 carbon atoms are highly mobile, colorless liquids with a characteristic pungent odor; acids with 4-9 carbon atoms are oily liquids with an unpleasant odor, and acids with a large number of carbon atoms are solids insoluble in water.

Chemical formulas of acids

Let us consider the chemical formulas of acids using the example of several representatives (both inorganic and organic): hydrochloric acid - HCl, sulfuric acid - H 2 SO 4, phosphoric acid - H 3 PO 4, acetic acid - CH 3 COOH and benzoic acid - C 6 H5COOH. The chemical formula shows the qualitative and quantitative composition of the molecule (how many and which atoms are included in a particular compound). Using the chemical formula, you can calculate the molecular weight of acids (Ar(H) = 1 amu, Ar(Cl) = 35.5 amu. amu, Ar(P) = 31 amu, Ar(O) = 16 amu, Ar(S) = 32 amu, Ar(C) = 12 a.m.):

Mr(HCl) = Ar(H) + Ar(Cl);

Mr(HCl) = 1 + 35.5 = 36.5.

Mr(H 2 SO 4) = 2×Ar(H) + Ar(S) + 4×Ar(O);

Mr(H 2 SO 4) = 2×1 + 32 + 4×16 = 2 + 32 + 64 = 98.

Mr(H 3 PO 4) = 3×Ar(H) + Ar(P) + 4×Ar(O);

Mr(H 3 PO 4) = 3×1 + 31 + 4×16 = 3 + 31 + 64 = 98.

Mr(CH 3 COOH) = 3×Ar(C) + 4×Ar(H) + 2×Ar(O);

Mr(CH 3 COOH) = 3×12 + 4×1 + 2×16 = 36 + 4 + 32 = 72.

Mr(C 6 H 5 COOH) = 7×Ar(C) + 6×Ar(H) + 2×Ar(O);

Mr(C 6 H 5 COOH) = 7 × 12 + 6 × 1 + 2 × 16 = 84 + 6 + 32 = 122.

Structural (graphic) formulas of acids

The structural (graphic) formula of a substance is more visual. It shows how atoms are connected to each other within a molecule. Let us indicate the structural formulas of each of the above compounds:

Rice. 1. Structural formula of hydrochloric acid.

Rice. 2. Structural formula of sulfuric acid.

Rice. 3. Structural formula of phosphoric acid.

Rice. 4. Structural formula of acetic acid.

Rice. 5. Structural formula of benzoic acid.

Ionic formulas

All inorganic acids are electrolytes, i.e. capable of dissociating in an aqueous solution into ions:

HCl ↔ H + + Cl - ;

H 2 SO 4 ↔ 2H + + SO 4 2- ;

H 3 PO 4 ↔ 3H + + PO 4 3- .

Examples of problem solving

EXAMPLE 1

EXAMPLE 2

Acids can be classified based on different criteria:

1) The presence of oxygen atoms in the acid

2) Acid basicity

The basicity of an acid is the number of “mobile” hydrogen atoms in its molecule, capable of being split off from the acid molecule in the form of hydrogen cations H + upon dissociation, 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 of 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.

Of the polybasic acids, the exception 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 products form a salt of the corresponding metal and hydrogen. 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 - a protective film of solid oxidation products is formed on the surface of such metals, 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.

Let's look at the most common acid formulas found in textbooks:

It is easy to notice that all acid formulas have in common the presence of hydrogen atoms (H), which comes first in the formula.

Determination of the valence of an acid residue

From the above list it can be seen that the number of these atoms may differ. Acids that contain only one hydrogen atom are called monobasic (nitric, hydrochloric, and others). Sulfuric, carbonic, and silicic acids are dibasic, since their formulas contain two H atoms. A tribasic phosphoric acid molecule contains three hydrogen atoms.

Thus, the amount of H in the formula characterizes the basicity of the acid.

The atom or group of atoms that are written after hydrogen are called acid residues. For example, in hydrosulfide acid the residue consists of one atom - S, and in phosphoric, sulfurous and many others - of two, and one of them is necessarily oxygen (O). On this basis, all acids are divided into oxygen-containing and oxygen-free.

Each acid residue has a certain valence. It is equal to the number of H atoms in the molecule of this acid. The valence of the HCl residue is equal to one, since it is a monobasic acid. Residues of nitric, perchloric, and nitrous acids have the same valency. The valency of the sulfuric acid residue (SO 4) is two, since there are two hydrogen atoms in its formula. Trivalent phosphoric acid residue.

Acidic residues - anions

In addition to valence, acid residues have charges and are anions. Their charges are indicated in the solubility table: CO 3 2−, S 2−, Cl− and so on. Please note: the charge of the acidic residue is numerically the same as its valency. For example, in silicic acid, the formula of which is H 2 SiO 3, the acid residue SiO 3 has a valence of II and a charge of 2-. Thus, knowing the charge of the acidic residue, it is easy to determine its valence and vice versa.

Summarize. Acids are compounds formed by hydrogen atoms and acidic residues. From the point of view of the theory of electrolytic dissociation, another definition can be given: acids are electrolytes, in solutions and melts of which hydrogen cations and anions of acid residues are present.

Hints

Chemical formulas of acids are usually learned by heart, as are their names. If you have forgotten how many hydrogen atoms are in a particular formula, but you know what its acidic residue looks like, the solubility table will come to your aid. The charge of the residue coincides in modulus with the valence, and that with the amount of H. For example, you remember that the remainder of carbonic acid is CO 3 . Using the solubility table, you determine that its charge is 2-, which means it is divalent, that is, carbonic acid has the formula H 2 CO 3.

There is often confusion with the formulas of sulfuric and sulfurous, as well as nitric and nitrous acids. Here, too, there is one point that makes it easier to remember: the name of the acid from the pair in which there are more oxygen atoms ends in -naya (sulfuric, nitric). An acid with fewer oxygen atoms in the formula has a name ending in -istaya (sulphurous, nitrogenous).

However, these tips will only help if the acid formulas are familiar to you. Let's repeat them again.