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 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 leaking chemical reactions acids with metals, a 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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| Acid | Acid residue | ||
| Formula | Name | Formula | Name |
| HBr | hydrobromic | Br – | bromide |
| HBrO3 | brominated | BrO3 – | bromate |
| HCN | hydrogen cyanide (cyanic) | CN- | cyanide |
| HCl | hydrochloric (hydrochloric) | Cl – | chloride |
| HClO | hypochlorous | ClO – | hypochlorite |
| HClO2 | chloride | ClO2 – | chlorite |
| HClO3 | hypochlorous | ClO3 – | chlorate |
| HClO4 | chlorine | ClO 4 – | perchlorate |
| H2CO3 | coal | HCO 3 – | bicarbonate |
| CO 3 2– | carbonate | ||
| H2C2O4 | sorrel | C2O42– | oxalate |
| CH3COOH | vinegar | CH 3 COO – | acetate |
| H2CrO4 | chrome | CrO 4 2– | chromate |
| H2Cr2O7 | dichrome | Cr 2 O 7 2– | dichromate |
| HF | hydrogen fluoride (fluoride) | F – | fluoride |
| HI | hydrogen iodide | I – | iodide |
| HIO 3 | iodic | IO 3 – | iodate |
| H2MnO4 | manganese | MnO 4 2– | manganate |
| HMnO4 | manganese | MnO4 – | permanganate |
| HNO2 | nitrogenous | NO 2 – | nitrite |
| HNO3 | nitrogen | NO 3 – | nitrate |
| H3PO3 | phosphorous | PO 3 3– | phosphite |
| H3PO4 | phosphorus | PO 4 3– | phosphate |
| HSCN | hydrothiocyanate (rhodanic) | SCN - | thiocyanate (rhodanide) |
| H2S | hydrogen sulfide | S 2– | sulfide |
| H2SO3 | sulfurous | SO 3 2– | sulfite |
| H2SO4 | sulfuric | SO 4 2– | sulfate |
End adj.
Prefixes most often used in names
Interpolation of reference values
Sometimes it is necessary to find out a density or concentration value that is not indicated in the reference tables. The required parameter can be found by interpolation.
Example
To prepare the HCl solution, the acid available in the laboratory was taken, the density of which was determined by a hydrometer. It turned out to be equal to 1.082 g/cm3.
According to the reference table, we find that an acid with a density of 1.080 has a mass fraction of 16.74%, and with 1.085 - 17.45%. To find the mass fraction of acid in an existing solution, we use the interpolation formula:
%,
where is the index 1 refers to a more dilute solution, and 2 - to more concentrated.
Preface……………………………..………….……….…......3
1. Basic concepts of titrimetric methods of analysis......7
2. Titration methods and methods……………………….....……...9
3. Calculation molar mass equivalents.…………………16
4. Methods of expressing the quantitative composition of solutions
in titrimetry……………………………………………………..21
4.1. Solution typical tasks on ways of expression
quantitative composition of solutions……………….……25
4.1.1. Calculation of the concentration of a solution based on the known mass and volume of the solution………………………………………………………..26
4.1.1.1. Problems for independent solution...29
4.1.2. Conversion of one concentration to another………...30
4.1.2.1. Problems for independent solution...34
5. Methods for preparing solutions…………………………...36
5.1. Solving typical problems for preparing solutions
in various ways…………………………………..39
5.2. Problems for independent solution………………….48
6. Calculation of titrimetric analysis results……….........51
6.1. Calculation of direct and substitution results
titration…………………………………………………………...51
6.2. Calculation of back titration results……………...56
7. Neutralization method (acid-base titration)……59
7.1. Examples of solving typical problems……………………..68
7.1.1. Direct and substitution titration……………68
7.1.1.1. Problems for independent solution...73
7.1.2. Back titration……………………………..76
7.1.2.1. Problems for independent solution...77
8. Oxidation-reduction method (redoximetry)………...80
8.1. Problems for independent solution………………….89
8.1.1. Redox reactions……..89
8.1.2. Calculation of titration results…………………...90
8.1.2.1. Substitution titration……………...90
8.1.2.2. Forward and reverse titration…………..92
9. Complexation method; complexometry…........94
9.1. Examples of solving typical problems……………………...102
9.2. Problems for independent solution………………...104
10. Deposition method……………………………………........106
10.1. Examples of solving typical problems…………………….110
10.2. Problems for independent solution……………….114
11. Individual tasks according to titrimetric
methods of analysis……………………………………………………………117
11.1. Plan for completing an individual task………...117
11.2. Options for individual tasks………………….123
Answers to problems………..……………………………………………………124
Symbols…………………………………………………….…127
Appendix……………………………………………………...128
EDUCATIONAL EDITION
ANALYTICAL CHEMISTRY
Acids- complex substances consisting of one or more hydrogen atoms that can be replaced by metal atoms and acidic residues.
Classification of acids
1. By the number of hydrogen atoms: number of hydrogen atoms ( n ) determines the basicity of acids:
n= 1 monobase
n= 2 dibase
n= 3 tribase
2. By composition:
a) Table of oxygen-containing acids, acid residues and corresponding acid oxides:
|
Acid (H n A) |
Acid residue (A) |
Corresponding acid oxide |
|
H 2 SO 4 sulfuric |
SO 4 (II) sulfate |
SO3 sulfur oxide (VI) |
|
HNO 3 nitrogen |
NO3(I)nitrate |
N 2 O 5 nitric oxide (V) |
|
HMnO 4 manganese |
MnO 4 (I) permanganate |
Mn2O7 manganese oxide ( VII) |
|
H 2 SO 3 sulfurous |
SO 3 (II) sulfite |
SO2 sulfur oxide (IV) |
|
H 3 PO 4 orthophosphoric |
PO 4 (III) orthophosphate |
P 2 O 5 phosphorus oxide (V) |
|
HNO 2 nitrogenous |
NO 2 (I) nitrite |
N 2 O 3 nitric oxide (III) |
|
H 2 CO 3 coal |
CO 3 (II) carbonate |
CO2 carbon monoxide ( IV) |
|
H 2 SiO 3 silicon |
SiO 3 (II) silicate |
SiO 2 silicon(IV) oxide |
|
HClO hypochlorous |
ClO(I) hypochlorite |
C l 2 O chlorine oxide (I) |
|
HClO 2 chloride |
ClO 2 (I) chlorite |
C l 2 O 3 chlorine oxide (III) |
|
HClO 3 chlorate |
ClO 3 (I) chlorate |
C l 2 O 5 chlorine oxide (V) |
|
HClO 4 chlorine |
ClO 4 (I) perchlorate |
C l 2 O 7 chlorine oxide (VII) |
b) Table of oxygen-free acids
|
Acid (H n A) |
Acid residue (A) |
|
HCl hydrochloric, hydrochloric |
Cl(I) chloride |
|
H 2 S hydrogen sulfide |
S(II) sulfide |
|
HBr hydrogen bromide |
Br(I) bromide |
|
HI hydrogen iodide |
I(I)iodide |
|
HF hydrogen fluoride, fluoride |
F(I) fluoride |
Physical properties of acids
Many acids, such as sulfuric, nitric, and hydrochloric, are colorless liquids. solid acids are also known: orthophosphoric, metaphosphoric HPO 3, boric H 3 BO 3 . Almost all acids are soluble in water. An example of an insoluble acid is silicic acid H2SiO3 . Acid solutions have a sour taste. For example, many fruits are given a sour taste by the acids they contain. Hence the names of acids: citric, malic, etc.
Methods for producing acids
|
oxygen-free |
oxygen-containing |
|
HCl, HBr, HI, HF, H2S |
HNO 3, H 2 SO 4 and others |
|
RECEIVING |
|
|
1. Direct interaction of nonmetals H 2 + Cl 2 = 2 HCl |
1. Acidic oxide + water = acid SO 3 + H 2 O = H 2 SO 4 |
|
2. Exchange reaction between salt and less volatile acid 2 NaCl (tv.) + H 2 SO 4 (conc.) = Na 2 SO 4 + 2HCl |
|
Chemical properties of acids
1. Change the color of the indicators
|
Indicator name |
Neutral environment |
Acidic environment |
|
Litmus |
Violet |
Red |
|
Phenolphthalein |
Colorless |
Colorless |
|
Methyl orange |
Orange |
Red |
|
Universal indicator paper |
Orange |
Red |
2. React with metals in the activity series up to H 2
(excl. HNO 3 -Nitric acid)
Video "Interaction of acids with metals"
Me + ACID = SALT + H 2 (r. substitution)
Zn + 2 HCl = ZnCl 2 + H 2
3. With basic (amphoteric) oxides – metal oxides
Video "Interaction of metal oxides with acids"
Fur x O y + ACID = SALT + H 2 O (exchange ruble)
4. React with bases – neutralization reaction
ACID + BASE= SALT+ H 2 O (exchange ruble)
H 3 PO 4 + 3 NaOH = Na 3 PO 4 + 3 H 2 O
5. React with salts of weak, volatile acids - if acid forms, precipitates or gas evolves:
2 NaCl (tv.) + H 2 SO 4 (conc.) = Na 2 SO 4 + 2HCl ( R . exchange )
Video "Interaction of acids with salts"
6. Decomposition of oxygen-containing acids when heated
(excl. H 2 SO 4 ; H 3 P.O. 4 )
ACID = ACID OXIDE + WATER (r. expansion)
Remember!Unstable acids (carbonic and sulfurous acids) - decompose into gas and water:
H 2 CO 3 ↔ H 2 O + CO 2
H 2 SO 3 ↔ H 2 O + SO 2
Hydrogen sulfide acid in products released as gas:
CaS + 2HCl = H 2 S+CaCl2
ASSIGNMENT TASKS
No. 1. Distribute chemical formulas acids in the table. Give them names:
LiOH, Mn 2 O 7, CaO, Na 3 PO 4, H 2 S, MnO, Fe (OH) 3, Cr 2 O 3, HI, HClO 4, HBr, CaCl 2, Na 2 O, HCl, H 2 SO 4, HNO 3, HMnO 4, Ca (OH) 2, SiO 2, Acids
Bes-sour-
native
Oxygen-containing
soluble
insoluble
one-
basic
two-basic
three-basic
No. 2. Write down the reaction equations:
Ca + HCl
Na+H2SO4
Al+H2S
Ca+H3PO4
Name the reaction products.
No. 3. Write down reaction equations and name the products:
Na 2 O + H 2 CO 3
ZnO + HCl
CaO + HNO3
Fe 2 O 3 + H 2 SO 4
No. 4. Write down equations for the reactions of acids with bases and salts:
KOH + HNO3
NaOH + H2SO3
Ca(OH) 2 + H 2 S
Al(OH) 3 + HF
HCl + Na 2 SiO 3
H2SO4 + K2CO3
HNO3 + CaCO3
Name the reaction products.
EXERCISES
Trainer No. 1. "Formula and names of acids"
Trainer No. 2. "Establishing correspondence: acid formula - oxide formula"
Safety precautions - First aid in case of acid contact with skin
Safety precautions -
| Acid formulas | Names of acids | Names of the corresponding salts |
| HClO4 | chlorine | perchlorates |
| HClO3 | hypochlorous | chlorates |
| HClO2 | chloride | chlorites |
| HClO | hypochlorous | hypochlorites |
| H5IO6 | iodine | periodates |
| HIO 3 | iodic | iodates |
| H2SO4 | sulfuric | sulfates |
| H2SO3 | sulfurous | sulfites |
| H2S2O3 | thiosulfur | thiosulfates |
| H2S4O6 | tetrathionic | tetrathionates |
| HNO3 | nitrogen | nitrates |
| HNO2 | nitrogenous | nitrites |
| H3PO4 | orthophosphoric | orthophosphates |
| HPO 3 | metaphosphoric | metaphosphates |
| H3PO3 | phosphorous | phosphites |
| H3PO2 | phosphorous | hypophosphites |
| H2CO3 | coal | carbonates |
| H2SiO3 | silicon | silicates |
| HMnO4 | manganese | permanganates |
| H2MnO4 | manganese | manganates |
| H2CrO4 | chrome | chromates |
| H2Cr2O7 | dichrome | dichromats |
| HF | hydrogen fluoride (fluoride) | fluorides |
| HCl | hydrochloric (hydrochloric) | chlorides |
| HBr | hydrobromic | bromides |
| HI | hydrogen iodide | iodides |
| H2S | hydrogen sulfide | sulfides |
| HCN | hydrogen cyanide | cyanides |
| HOCN | cyan | cyanates |
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 the remainder of hydrochloric acid(Cl). Name of salt: iron (III) chloride. Please note: in in this case we must not only name the metal, but also indicate its valency (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.
If you are interested not only in the nomenclature of salts, but also in the methods of their preparation and Chemical properties, I recommend turning to the relevant sections of the chemistry reference book: "
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
| Acid formula | Acid name | Acid residue (anion) | Name of salts (average) |
|---|---|---|---|
| HF | Hydrofluoric (fluoric) | F − | Fluorides |
| HCl | Hydrochloric (hydrochloric) | Cl − | Chlorides |
| HBr | Hydrobromic | Br− | Bromides |
| HI | Hydroiodide | I − | Iodides |
| H2S | Hydrogen sulfide | S 2− | Sulfides |
| H2SO3 | Sulphurous | SO 3 2 − | Sulfites |
| H2SO4 | Sulfuric | SO 4 2 − | Sulfates |
| HNO2 | Nitrogenous | NO2− | Nitrites |
| HNO3 | Nitrogen | NO 3 − | Nitrates |
| H2SiO3 | Silicon | SiO 3 2 − | Silicates |
| HPO 3 | Metaphosphoric | PO 3 − | Metaphosphates |
| H3PO4 | Orthophosphoric | PO 4 3 − | Orthophosphates (phosphates) |
| H4P2O7 | Pyrophosphoric (biphosphoric) | P 2 O 7 4 − | Pyrophosphates (diphosphates) |
| HMnO4 | Manganese | MnO 4 − | Permanganates |
| H2CrO4 | Chrome | CrO 4 2 − | Chromates |
| H2Cr2O7 | Dichrome | Cr 2 O 7 2 − | Dichromates (bichromates) |
| H2SeO4 | Selenium | SeO 4 2 − | Selenates |
| H3BO3 | Bornaya | BO 3 3 − | Orthoborates |
| HClO | Hypochlorous | ClO – | Hypochlorites |
| HClO2 | Chloride | ClO2− | Chlorites |
| HClO3 | Chlorous | ClO3− | Chlorates |
| HClO4 | Chlorine | ClO 4 − | Perchlorates |
| H2CO3 | Coal | CO 3 3 − | Carbonates |
| CH3COOH | Vinegar | CH 3 COO − | Acetates |
| HCOOH | Ant | HCOO − | Formiates |
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 “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
| Classification sign | Acid type | Examples |
|---|---|---|
| Number of hydrogen ions formed upon complete dissociation of an acid molecule | Monobase | HCl, HNO3, CH3COOH |
| Dibasic | H2SO4, H2S, H2CO3 | |
| Tribasic | H3PO4, H3AsO4 | |
| The presence or absence of an oxygen atom in a molecule | Oxygen-containing (acid hydroxides, oxoacids) | HNO2, H2SiO3, H2SO4 |
| Oxygen-free | HF, H2S, HCN | |
| Degree of dissociation (strength) | Strong (completely dissociate, strong electrolytes) | HCl, HBr, HI, H 2 SO 4 (diluted), HNO 3, HClO 3, HClO 4, HMnO 4, H 2 Cr 2 O 7 |
| Weak (partially dissociate, weak electrolytes) | HF, HNO 2, H 2 SO 3, HCOOH, CH 3 COOH, H 2 SiO 3, H 2 S, HCN, H 3 PO 4, H 3 PO 3, HClO, HClO 2, H 2 CO 3, H 3 BO 3, H 2 SO 4 (conc) | |
| Oxidative properties | Oxidizing agents due to H + ions (conditionally non-oxidizing acids) | HCl, HBr, HI, HF, H 2 SO 4 (dil), H 3 PO 4, CH 3 COOH |
| Oxidizing agents due to anion (oxidizing acids) | HNO 3, HMnO 4, H 2 SO 4 (conc), H 2 Cr 2 O 7 | |
| Anion reducing agents | HCl, HBr, HI, H 2 S (but not HF) | |
| Thermal stability | Exist only in solutions | H 2 CO 3, H 2 SO 3, HClO, HClO 2 |
| Easily decomposes when heated | H 2 SO 3 , HNO 3 , H 2 SiO 3 | |
| Thermally stable | H 2 SO 4 (conc), H 3 PO 4 |
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 ↓