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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(H2SO4 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 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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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 a theoretical point of view 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.

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. With the help of characteristic reactions, the presence of certain ions in solutions and in the precipitate, which must first be dissolved, is proven.

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.

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, as in individual form, and in solutions H 2 SO 4 , HNO 3 , H 3 PO 4 , CH 3 COOH are known.

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 “solution of hydrochloric acid" wrong.

Most acids are soluble in water; silicic acid H 2 SiO 3 is insoluble. The overwhelming majority of acids have a molecular structure. Examples structural formulas 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 acids are caused by 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, the litmus is not red, but pink, the solution above the precipitate is 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:

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

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);

BaO + SO 2 = BaSO 3

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).

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 Anoxic acids

• Oxygen-containing acids

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):

AgNO 3 + HCl = AgCl↓ + HNO 3

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

Let me briefly remind you of specific examples how to properly call salts.

Example 1. The salt K 2 SO 4 is formed by a sulfuric acid residue (SO 4) and metal K. Salts of sulfuric acid are called sulfates. K 2 SO 4 - potassium sulfate.

Example 2. FeCl 3 - the salt contains iron and a hydrochloric acid residue (Cl). Name of salt: iron (III) chloride. Please note: in this case we must not only name the metal, but also indicate its valence (III). In the previous example, this was not necessary, since the valency of sodium is constant.

Important: the name of the salt should indicate the valence of the metal only if the metal has a variable valency!

Example 3. Ba(ClO) 2 - the salt contains barium and the remainder of hypochlorous acid (ClO). Salt name: barium hypochlorite. The valency of the metal Ba in all its compounds is two; it does not need to be indicated.

Example 4. (NH 4) 2 Cr 2 O 7. The NH 4 group is called ammonium, the valence of this group is constant. Name of salt: ammonium dichromate (dichromate).

In the above examples we only encountered the so-called. medium or normal salts. Acidic, basic, double and complex salts, salts of organic acids will not be discussed here.