The reaction of acid salt formation. Salts

Every day we come across salts and don’t even think about the role they play in our lives. But without them, the water would not be so tasty, and the food would not bring pleasure, and the plants would not grow, and life on earth could not exist if there were no salt in our world. So what are these substances and what properties of salts make them irreplaceable?

What are salts

In terms of its composition, this is the most numerous class, characterized by diversity. Back in the 19th century, the chemist J. Werzelius defined a salt as a product of a reaction between an acid and a base, in which a hydrogen atom is replaced by a metal one. In water, salts usually dissociate into a metal or ammonium (cation) and an acidic residue (anion).

You can get salts in the following ways:

• through the interaction of a metal and a non-metal, in this case it will be oxygen-free;
• when a metal reacts with an acid, a salt is obtained and hydrogen is released;
• a metal can displace another metal from solution;
• when two oxides interact - acidic and basic (they are also called non-metal oxide and metal oxide, respectively);
• the reaction of a metal oxide and an acid produces salt and water;
• the reaction between a base and a nonmetal oxide also produces salt and water;
• using an ion exchange reaction, in this case various water-soluble substances (bases, acids, salts) can react, but the reaction will proceed if gas, water or slightly soluble (insoluble) salts are formed in water.

The properties of salts depend only on the chemical composition. But first, let's look at their classes.

Classification

Depending on the composition, the following classes of salts are distinguished:

• by oxygen content (oxygen-containing and oxygen-free);
• by interaction with water (soluble, slightly soluble and insoluble).

This classification does not fully reflect the diversity of substances. The modern and most complete classification, reflecting not only the composition, but also the properties of salts, is presented in the following table.

Physical properties

No matter how wide the class of these substances is, it is possible to identify the general physical properties of salts. These are substances of non-molecular structure, with an ionic crystal lattice.

Very high melting and boiling points. Under normal conditions, all salts do not conduct electricity, but in solution, most of them conduct electricity perfectly.

The color can be very different, it depends on the metal ion included in its composition. Ferrous sulfate (FeSO 4) is green, ferrous chloride (FeCl 3) is dark red, and potassium chromate (K 2 CrO 4) is a beautiful bright yellow color. But most salts are still colorless or white.

Solubility in water also varies and depends on the composition of the ions. In principle, all physical properties of salts have a peculiarity. They depend on which metal ion and which acid residue are included in the composition. Let's continue looking at salts.

Chemical properties of salts

There is also an important feature here. Like the physical, chemical properties of salts depend on their composition. And also on what class they belong to.

But the general properties of salts can still be highlighted:

• many of them decompose when heated to form two oxides: acidic and basic, and oxygen-free ones - metal and non-metal;
• salts also interact with other acids, but the reaction occurs only if the salt contains an acidic residue of a weak or volatile acid or the result is an insoluble salt;
• interaction with alkali is possible if the cation forms an insoluble base;
• a reaction between two different salts is also possible, but only if one of the newly formed salts does not dissolve in water;
• A reaction with a metal can also occur, but it is only possible if we take a metal located to the right in the voltage series from the metal contained in the salt.

The chemical properties of salts classified as normal are discussed above, but other classes react with substances somewhat differently. But the difference is only in the output products. Basically, all the chemical properties of the salts are preserved, as are the requirements for the reactions.

Chemical properties of salts

Salts should be considered as the product of the reaction of an acid and a base. As a result, the following may form:

1. normal (average) - are formed when the amount of acid and base is sufficient for complete interaction. Names of normal salts They consist of two parts. First the anion (acid residue) is called, then the cation.
2. sour - are formed when there is an excess of acid and an insufficient amount of alkali, because in this case there are not enough metal cations to replace all the hydrogen cations present in the acid molecule. You will always see hydrogen in the acidic residues of this type of salt. Acid salts are formed only by polybasic acids and exhibit the properties of both salts and acids. In the names of acid salts a prefix is ​​placed hydro- to the anion.
3. basic salts - are formed when there is an excess of base and an insufficient amount of acid, because in this case the anions of acidic residues are not enough to completely replace the hydroxyl groups present in the base. the main salts in the cations contain hydroxo groups. Basic salts are possible for polyacid bases, but not for monoacid bases. Some basic salts are capable of decomposing independently, releasing water in the process, forming oxo salts that have the properties of basic salts. Name of main salts is constructed as follows: a prefix is ​​added to the anion hydroxo-.

Typical reactions of normal salts

• They react well with metals. At the same time, more active metals displace less active ones from solutions of their salts.
• With acids, alkalis and other salts, reactions proceed to completion, provided that a precipitate, gas or poorly dissociable compounds are formed.
• In the reactions of salts with alkalis, substances such as nickel (II) hydroxide Ni(OH) 2 are formed - a precipitate; ammonia NH 3 – gas; water H 2 O is a weak electrolyte, a poorly dissociated compound:
• Salts react with each other if a precipitate is formed or if a more stable compound is formed.
• Many normal salts decompose when heated to form two oxides - acidic and basic.
• Nitrates decompose in a different way from other normal salts. When heated, nitrates of alkali and alkaline earth metals release oxygen and turn into nitrites:
• Nitrates of almost all other metals decompose to oxides:
• Nitrates of some heavy metals (silver, mercury, etc.) decompose when heated to metals:

Typical reactions of acid salts

• They enter into all the reactions that acids enter into. They react with alkalis; if the acid salt and alkali contain the same metal, then a normal salt is formed as a result.
• If the alkali contains another metal, then double salts are formed.

Typical reactions of basic salts

• These salts undergo the same reactions as bases. They react with acids; if the basic salt and the acid contain the same acidic residue, then the result is a normal salt.
• If the acid contains another acid residue, then double salts are formed.

Complex salts- a compound whose crystal lattice sites contain complex ions.

Video tutorial 1: Classification of inorganic salts and their nomenclature

Video tutorial 2: Methods for obtaining inorganic salts. Chemical properties of salts

Lecture: Characteristic chemical properties of salts: medium, acidic, basic; complex (using the example of aluminum and zinc compounds)

Salts- these are chemical compounds consisting of metal cations (or ammonium) and acidic residues.

Salts should also be considered as a product of the interaction of an acid and a base. As a result of this interaction, the following can form:

normal (average),

• basic salts.

Normal salts are formed when the amount of acid and base is sufficient for complete interaction. Eg:

H 3 PO 4 + 3KON → K 3 PO 4 + 3H 2 O.

The names of normal salts consist of two parts. First the anion (acid residue) is called, then the cation. For example: sodium chloride - NaCl, iron(III) sulfate - Fe 2 (SO 4) 3, potassium carbonate - K 2 CO 3, potassium phosphate - K 3 PO 4, etc.

Acid salts are formed when there is an excess of acid and an insufficient amount of alkali, because in this case there are not enough metal cations to replace all the hydrogen cations present in the acid molecule. Eg:

H 3 PO 4 + 2KON = K 2 NPO 4 + 2H 2 O;

H 3 PO 4 + KOH = KH 2 PO 4 + H 2 O.

You will always see hydrogen in the acidic residues of this type of salt. Acid salts are always possible for polybasic acids, but not for monobasic acids.

The names of acidic salts are prefixed hydro- to the anion. For example: iron(III) hydrogen sulfate - Fe(HSO 4) 3, potassium hydrogen carbonate - KHCO 3, potassium hydrogen phosphate - K 2 HPO 4, etc.

Basic salts are formed when there is an excess of base and an insufficient amount of acid, because in this case the anions of acidic residues are not enough to completely replace the hydroxyl groups present in the base. Eg:

Cr(OH) 3 + HNO 3 → Cr(OH) 2 NO 3 + H 2 O;

Cr(OH) 3 + 2HNO 3 → CrOH(NO 3) 2 + 2H 2 O.

Thus, the main salts in the cations contain hydroxo groups. Basic salts are possible for polyacid bases, but not for monoacid bases. Some basic salts are capable of decomposing independently, releasing water in the process, forming oxo salts that have the properties of basic salts. Eg:

Sb(OH) 2 Cl → SbOCl + H 2 O;

Bi(OH) 2 NO 3 → BiONO 3 + H 2 O.

The name of the main salts is constructed as follows: the prefix is ​​added to the anion hydroxo-. For example: iron(III) hydroxy sulfate - FeOHSO 4, aluminum hydroxy sulfate - AlOHSO 4, iron (III) dihydroxochloride - Fe(OH) 2 Cl, etc.

Many salts, being in a solid state of aggregation, are crystalline hydrates: CuSO4.5H2O; Na2CO3.10H2O, etc.

Chemical properties of salts

Salts are fairly solid crystalline substances that have ionic bonds between cations and anions. The properties of salts are determined by their interaction with metals, acids, bases and salts.

Typical reactions of normal salts

They react well with metals. At the same time, more active metals displace less active ones from solutions of their salts. Eg:

Zn + CuSO 4 → ZnSO 4 + Cu;

Cu + Ag 2 SO 4 → CuSO 4 + 2Ag.

With acids, alkalis and other salts, reactions proceed to completion, provided that a precipitate, gas or poorly dissociable compounds are formed. For example, in reactions of salts with acids, substances such as hydrogen sulfide H 2 S are formed - gas; barium sulfate BaSO 4 – sediment; acetic acid CH 3 COOH is a weak electrolyte, a poorly dissociated compound. Here are the equations for these reactions:

K 2 S + H 2 SO 4 → K 2 SO 4 + H 2 S;

BaCl 2 + H 2 SO 4 → BaSO 4 + 2HCl;

CH 3 COONa + HCl → NaCl + CH 3 COOH.

In the reactions of salts with alkalis, substances such as nickel (II) hydroxide Ni(OH) 2 are formed - a precipitate; ammonia NH 3 – gas; water H 2 O is a weak electrolyte, a poorly dissociated compound:

NiCl 2 + 2KOH → Ni(OH) 2 + 2KCl;

NH 4 Cl + NaOH → NH 3 +H 2 O +NaCl.

Salts react with each other if a precipitate forms:

Ca(NO 3) 2 + Na 2 CO 3 → 2NaNO 3 + CaCO 3.

Or in the case of a more stable connection:

Ag 2 CrO 4 + Na 2 S → Ag 2 S + Na 2 CrO 4.

In this reaction, black silver sulfide is formed from brick-red silver chromate, due to the fact that it is a more insoluble precipitate than chromate.

Many normal salts decompose when heated to form two oxides - acidic and basic:

CaCO 3 → CaO + CO 2.

Nitrates decompose in a different way from other normal salts. When heated, nitrates of alkali and alkaline earth metals release oxygen and turn into nitrites:

2NaNO 3 → 2NaNO 2 + O 2.

Nitrates of almost all other metals decompose to oxides:

2Zn(NO 3) 2 → 2ZnO + 4NO 2 + O 2.

Nitrates of some heavy metals (silver, mercury, etc.) decompose when heated to metals:

2AgNO3 → 2Ag + 2NO2 + O2.

A special position is occupied by ammonium nitrate, which, up to the melting point (170 o C), partially decomposes according to the equation:

NH 4 NO 3 → NH 3 + HNO 3 .

At temperatures 170 - 230 o C, according to the equation:

NH 4 NO 3 → N 2 O + 2H 2 O.

At temperatures above 230 o C - with an explosion, according to the equation:

2NH 4 NO 3 → 2N 2 + O 2 + 4H 2 O.

Ammonium chloride NH 4 Cl decomposes to form ammonia and hydrogen chloride:

NH 4 Cl → NH 3 + HCl.

Typical reactions of acid salts

NaH CO3+ Na OH→ Na 2 CO3+ H 2 O .

NaHCO 3 +Li OH → Li NaCO 3+ H 2 O .

Typical reactions main salts

Cu( OH)Cl+ H Cl → Cu Cl 2 + H 2 O .

Cu( OH)Cl + HBr → Cu Br Cl+ H 2 O .

Complex salts

Complex connection- a compound whose crystal lattice sites contain complex ions.

Let's consider complex compounds of aluminum - tetrahydroxoaluminates and zinc - tetrahydroxoaluminates. Complex ions are indicated in square brackets in the formulas of these substances.

Chemical properties of sodium tetrahydroxoaluminate Na and sodium tetrahydroxoaluminate Na 2:

1. Like all complex compounds, the above substances dissociate:

• Na → Na + + - ;
• Na 2 → 2Na + + - .

Please note that further dissociation of complex ions is not possible.

2. In reactions with excess strong acids, two salts are formed. Consider the reaction of sodium tetrahydroxoaluminate with a dilute solution of hydrogen chloride:

• Na + 4HCl→ Al Cl 3 + Na Cl + H2O.

We see the formation of two salts: aluminum chloride, sodium chloride and water. A similar reaction will occur in the case of sodium tetrahydroxycinate.

3. If a strong acid is not enough, let's say instead 4 HCl We took 2 HCl, then the salt forms the most active metal, in this case sodium is more active, which means sodium chloride is formed, and the resulting hydroxides of aluminum and zinc will precipitate. Let us consider this case using the reaction equation with sodium tetrahydroxycinate:

Na 2 + 2HCl→ 2Na Cl+ Zn (OH) 2 ↓ +2H2O.

Reasons

Bases are compounds containing only hydroxide ions OH - as an anion. The number of hydroxide ions that can be replaced by an acidic residue determines the acidity of the base. In this regard, bases are one-, two- and polyacid; however, true bases most often include one- and two-acid. Among them, water-soluble and water-insoluble bases should be distinguished. Please note that bases that are soluble in water and dissociate almost completely are called alkalis (strong electrolytes). These include hydroxides of alkali and alkaline earth elements and in no case a solution of ammonia in water.

The name of the base begins with the word hydroxide, after which the Russian name of the cation is given in the genitive case, and its charge is indicated in parentheses. It is allowed to list the number of hydroxide ions using the prefixes di-, tri-, tetra. For example: Mn(OH) 3 - manganese (III) hydroxide or manganese trihydroxide.

Note that there is a genetic relationship between bases and basic oxides: basic oxides correspond to bases. Therefore, base cations most often have a charge of one or two, which corresponds to the lowest oxidation states of metals.

Remember the basic ways to obtain bases

1. Interaction of active metals with water:

2Na + 2H 2 O = 2NaOH + H 2

La + 6H 2 O = 2La(OH) 3 + 3H 2

Interaction of basic oxides with water:

CaO + H 2 O = Ca (OH) 2

MgO + H 2 O = Mg(OH) 2.

3. Interaction of salts with alkalis:

MnSO 4 + 2KOH = Mn(OH) 2 ↓ + K 2 SO 4

NH 4 С1 + NaOH = NaCl + NH 3 ∙ H 2 O

Na 2 CO 3 + Ca(OH) 2 = 2NaOH + CaCO 3

MgOHCl + NaOH = Mg(OH) 2 + NaCl.

Electrolysis of aqueous salt solutions with a diaphragm:

2NaCl + 2H 2 O → 2NaOH + Cl 2 + H 2

Please note that in step 3, the starting reagents must be selected in such a way that among the reaction products there is either a sparingly soluble compound or a weak electrolyte.

Note that when considering the chemical properties of bases, reaction conditions depend on the solubility of the base.

1. Interaction with acids:

NaOH + H 2 SO 4 = NaHSO 4 + H 2 O

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

2Mg(OH) 2 + H 2 SO 4 = (MgOH) 2 SO 4 + 2H 2 O

Mg(OH) 2 + H 2 SO 4 = MgSO 4 + 2H 2 O

Mg(OH) 2 + 2H 2 SO 4 = Mg(HSO 4) 2 + 2H 2 O

2. Interaction with acid oxides:

NaOH + CO 2 = NaHCO 3

2NaOH + CO 2 = Na 2 CO 3 + H 2 O

Fe(OH) 2 + P 2 O 5 = Fe(PO 3) 2 + H 2 O

3Fe(OH) 2 + P 2 O 5 = Fe 3 (PO 4) 2 + 2H 2 O

3. Interaction with amphoteric oxides:

A1 2 O 3 + 2NaOH p + 3H 2 O = 2Na

Al 2 O 3 + 2NaOH T = 2NaAlO 2 + H 2 O

Cr 2 O 3 + Mg(OH) 2 = Mg(CrO 2) 2 + H 2 O

4. Interaction with ampheteric hydroxides:

Ca(OH) 2 + 2Al(OH) 3 = Ca(AlO 2) 2 + 4H 2 O

3NaOH + Cr(OH) 3 = Na 3

Interaction with salts.

To the reactions described in point 3 of the production methods, the following should be added:

2ZnSO 4 + 2KOH = (ZnOH) 2 S0 4 + K 2 SO 4

NaHCO 3 + NaOH = Na 2 CO 3 + H 2 O

BeSO 4 + 4NaOH = Na 2 + Na 2 SO 4

Cu(OH) 2 + 4NH 3 ∙H 2 O = (OH) 2 + 4H 2 O

6. Oxidation to amphoteric hydroxides or salts:

4Fe(OH) 2 + O 2 + 2H 2 O = 4Fe(OH) 3

2Cr(OH) 2 + 2H 2 O + Na 2 O 2 + 4NaOH = 2Na 3.

7. Heat decomposition:

Ca(OH) 2 = CaO + H 2 O.

Please note that alkali metal hydroxides, except lithium, do not participate in such reactions.

!!!Are there alkaline precipitations?!!! Yes, there are, but they are not as widespread as acid precipitation, are little known, and their impact on environmental objects has been practically unstudied. Nevertheless, their consideration deserves attention.

The origin of alkaline precipitation can be explained as follows.

CaCO 3 →CaO + CO 2

In the atmosphere, calcium oxide combines with water vapor during condensation, with rain or sleet, forming calcium hydroxide:

CaO + H 2 O →Ca(OH) 2,

which creates an alkaline reaction of atmospheric precipitation. In the future, it is possible to react calcium hydroxide with carbon dioxide and water to form calcium carbonate and calcium bicarbonate:

Ca(OH) 2 + CO 2 → CaCO 3 + H 2 O;

CaCO 3 + CO 2 + H 2 O → Ca(HC0 3) 2.

Chemical analysis of rainwater showed that it contains sulfate and nitrate ions in small quantities (about 0.2 mg/l). As is known, the cause of the acidic nature of precipitation is sulfuric and nitric acids. At the same time, there is a high content of calcium cations (5-8 mg/l) and bicarbonate ions, the content of which in the area of ​​​​the construction complex enterprises is 1.5-2 times higher than in other areas of the city, and amounts to 18-24 mg /l. This shows that the calcium carbonate system and the processes occurring in it play a major role in the formation of local alkaline sediments, as mentioned above.

Alkaline precipitation affects plants; changes in the phenotypic structure of plants are noted. There are traces of “burns” on the leaf blades, a white coating on the leaves and a depressed state of herbaceous plants.

1. Bases react with acids to form salt and water:

Cu(OH) 2 + 2HCl = CuCl 2 + 2H 2 O

2. With acid oxides, forming salt and water:

Ca(OH) 2 + CO 2 = CaCO 3 + H 2 O

3. Alkalis react with amphoteric oxides and hydroxides, forming salt and water:

2NaOH + Cr 2 O 3 = 2NaCrO 2 + H 2 O

KOH + Cr(OH) 3 = KCrO 2 + 2H 2 O

4. Alkalis react with soluble salts, forming either a weak base, a precipitate, or a gas:

2NaOH + NiCl 2 = Ni(OH) 2 ¯ + 2NaCl

base

2KOH + (NH 4) 2 SO 4 = 2NH 3 + 2H 2 O + K 2 SO 4

Ba(OH) 2 + Na 2 CO 3 = BaCO 3 ¯ + 2NaOH

5. Alkalis react with some metals, which correspond to amphoteric oxides:

2NaOH + 2Al + 6H 2 O = 2Na + 3H 2

6. Effect of alkali on the indicator:

OH - + phenolphthalein ® crimson color

OH - + litmus ® blue color

7. Decomposition of some bases when heated:

Сu(OH) 2 ® CuO + H 2 O

Amphoteric hydroxides– chemical compounds exhibiting the properties of both bases and acids. Amphoteric hydroxides correspond to amphoteric oxides (see paragraph 3.1).

Amphoteric hydroxides are usually written in the form of a base, but they can also be represented in the form of an acid:

Zn(OH) 2 Û H 2 ZnO 2

foundation

Chemical properties of amphoteric hydroxides

1. Amphoteric hydroxides interact with acids and acid oxides:

Be(OH) 2 + 2HCl = BeCl 2 + 2H 2 O

Be(OH) 2 + SO 3 = BeSO 4 + H 2 O

2. Interact with alkalis and basic oxides of alkali and alkaline earth metals:

Al(OH) 3 + NaOH = NaAlO 2 + 2H 2 O;

H 3 AlO 3 acid sodium metaaluminate

(H 3 AlO 3 ® HAlO 2 + H 2 O)

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

All amphoteric hydroxides are weak electrolytes

Salts

Salts- These are complex substances consisting of metal ions and an acid residue. Salts are products of complete or partial replacement of hydrogen ions with metal (or ammonium) ions in acids. Types of salts: medium (normal), acidic and basic.

Medium salts- these are the products of complete replacement of hydrogen cations in acids with metal (or ammonium) ions: Na 2 CO 3, NiSO 4, NH 4 Cl, etc.

Chemical properties of medium salts

1. Salts interact with acids, alkalis and other salts, forming either a weak electrolyte or a precipitate; or gas:

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

Na 2 SO 4 + Ba(OH) 2 = BaSO 4 ¯ + 2NaOH

CaCl 2 + 2AgNO 3 = 2AgCl¯ + Ca(NO 3) 2

2CH 3 COONa + H 2 SO 4 = Na 2 SO 4 + 2CH 3 COOH

NiSO 4 + 2KOH = Ni(OH) 2 ¯ + K 2 SO 4

base

NH 4 NO 3 + NaOH = NH 3 + H 2 O + NaNO 3

2. Salts interact with more active metals. A more active metal displaces a less active metal from the salt solution (Appendix 3).

Zn + CuSO 4 = ZnSO 4 + Cu

Acid salts- these are products of incomplete replacement of hydrogen cations in acids with metal (or ammonium) ions: NaHCO 3, NaH 2 PO 4, Na 2 HPO 4, etc. Acid salts can only be formed by polybasic acids. Almost all acid salts are highly soluble in water.

Obtaining acidic salts and converting them to medium salts

1. Acid salts are obtained by reacting an excess of acid or acid oxide with a base:

H 2 CO 3 + NaOH = NaHCO 3 + H 2 O

CO 2 + NaOH = NaHCO 3

2. When excess acid interacts with the basic oxide:

2H 2 CO 3 + CaO = Ca(HCO 3) 2 + H 2 O

3. Acid salts are obtained from medium salts by adding acid:

· eponymous

Na 2 SO 3 + H 2 SO 3 = 2NaHSO 3;

Na 2 SO 3 + HCl = NaHSO 3 + NaCl

4. Acid salts are converted to medium salts using alkali:

NaHCO 3 + NaOH = Na 2 CO 3 + H 2 O

Basic salts– these are products of incomplete substitution of hydroxo groups (OH - ) bases with an acidic residue: MgOHCl, AlOHSO 4, etc. Basic salts can only be formed by weak bases of polyvalent metals. These salts are generally sparingly soluble.

Obtaining basic salts and converting them to medium salts

1. Basic salts are obtained by reacting an excess of base with an acid or acid oxide:

Mg(OH) 2 + HCl = MgOHCl¯ + H 2 O

hydroxo-

magnesium chloride

Fe(OH) 3 + SO 3 = FeOHSO 4 ¯ + H 2 O

hydroxo-

iron(III) sulfate

2. Basic salts are formed from medium salt by adding a lack of alkali:

Fe 2 (SO 4) 3 + 2NaOH = 2FeOHSO 4 + Na 2 SO 4

3. Basic salts are converted to medium salts by adding an acid (preferably the one that corresponds to the salt):

MgOHCl + HCl = MgCl 2 + H 2 O

2MgOHCl + H 2 SO 4 = MgCl 2 + MgSO 4 + 2H 2 O

ELECTROLYTES

Electrolytes- these are substances that disintegrate into ions in solution under the influence of polar solvent molecules (H 2 O). Based on their ability to dissociate (break down into ions), electrolytes are conventionally divided into strong and weak. Strong electrolytes dissociate almost completely (in dilute solutions), while weak electrolytes dissociate into ions only partially.

Strong electrolytes include:

· strong acids (see p. 20);

· strong bases – alkalis (see p. 22);

· almost all soluble salts.

Weak electrolytes include:

weak acids (see p. 20);

· bases are not alkali;

One of the main characteristics of a weak electrolyte is dissociation constant – TO . For example, for a monobasic acid,

HA Û H + +A - ,

where, is the equilibrium concentration of H + ions;

– equilibrium concentration of acid anions A - ;

– equilibrium concentration of acid molecules,

Or for a weak foundation,

MOH Û M + +OH - ,

,

where, is the equilibrium concentration of M + cations;

– equilibrium concentration of hydroxide ions OH - ;

– equilibrium concentration of weak base molecules.

Dissociation constants of some weak electrolytes (at t = 25°C)