Weak electrolytes

Weak electrolytes Substances that partially dissociate into ions. Solutions of weak electrolytes, along with ions, contain undissociated molecules. Weak electrolytes cannot give a high concentration of ions in solution. Weak electrolytes include:

1) almost all organic acids (CH 3 COOH, C 2 H 5 COOH, etc.);

2) some inorganic acids (H 2 CO 3 , H 2 S, etc.);

3) almost all water-soluble salts, bases and ammonium hydroxide Ca 3 (PO 4) 2 ; Cu(OH) 2 ; Al(OH) 3 ; NH4OH;

They are poor conductors (or almost non-conductors) of electricity.

Ion concentrations in solutions of weak electrolytes are qualitatively characterized by the degree and dissociation constant.

The degree of dissociation is expressed in fractions of a unit or as a percentage (a \u003d 0.3 is the conditional division boundary into strong and weak electrolytes).

The degree of dissociation depends on the concentration of the weak electrolyte solution. When diluted with water, the degree of dissociation always increases, because the number of solvent molecules (H 2 O) increases per solute molecule. According to the Le Chatelier principle, the equilibrium of electrolytic dissociation in this case should shift in the direction of product formation, i.e. hydrated ions.

The degree of electrolytic dissociation depends on the temperature of the solution. Usually, with increasing temperature, the degree of dissociation increases, because bonds in molecules are activated, they become more mobile and easier to ionize. The concentration of ions in a weak electrolyte solution can be calculated knowing the degree of dissociation a and the initial concentration of the substance c in solution.

HAn = H + + An - .

The equilibrium constant K p of this reaction is the dissociation constant K d:

K d = . / . (10.11)

If we express the equilibrium concentrations in terms of the concentration of a weak electrolyte C and its degree of dissociation α, then we get:

K d \u003d C. α. C. α/C. (1-α) = C. α 2 /1-α. (10.12)

This relationship is called Ostwald's dilution law. For very weak electrolytes at α<<1 это уравнение упрощается:

K d \u003d C. α 2. (10.13)

This allows us to conclude that, at infinite dilution, the degree of dissociation α tends to unity.

Protolytic equilibrium in water:

,

,

At a constant temperature in dilute solutions, the concentration of water in water is constant and equal to 55.5, ( )

, (10.15)

where K in is the ionic product of water.

Then =10 -7 . In practice, due to the convenience of measuring and recording, a value is used - the pH value, (criterion) of the strength of an acid or base. Similarly .

From equation (11.15): . At pH = 7 - the reaction of the solution is neutral, at pH<7 – кислая, а при pH>7 - alkaline.

Under normal conditions (0°C):

, then

Figure 10.4 - pH of various substances and systems

10.7 Solutions of strong electrolytes

Strong electrolytes are substances that, when dissolved in water, almost completely decompose into ions. As a rule, strong electrolytes include substances with ionic or highly polar bonds: all highly soluble salts, strong acids (HCl, HBr, HI, HClO 4, H 2 SO 4, HNO 3) and strong bases (LiOH, NaOH, KOH, RbOH, CsOH, Ba (OH) 2, Sr (OH) 2, Ca (OH) 2).

In a solution of a strong electrolyte, the solute is found mainly in the form of ions (cations and anions); undissociated molecules are practically absent.

The fundamental difference between strong and weak electrolytes is that the dissociation equilibrium of strong electrolytes is completely shifted to the right:

H 2 SO 4 \u003d H + + HSO 4 -,

and therefore the constant of equilibrium (dissociation) turns out to be an indeterminate quantity. The decrease in electrical conductivity with increasing concentration of a strong electrolyte is due to the electrostatic interaction of ions.

The Dutch scientist Petrus Josephus Wilhelmus Debye and the German scientist Erich Hückel postulated:

1) the electrolyte completely dissociates, but in relatively dilute solutions (C M = 0.01 mol. l -1);

2) each ion is surrounded by a shell of ions of the opposite sign. In turn, each of these ions is solvated. This environment is called the ionic atmosphere. In the electrolytic interaction of ions of opposite signs, it is necessary to take into account the influence of the ionic atmosphere. When a cation moves in an electrostatic field, the ionic atmosphere is deformed; it thickens before him and thins behind him. This asymmetry of the ionic atmosphere has the more inhibitory effect on the movement of the cation, the higher the concentration of electrolytes and the greater the charge of the ions. In these systems, the concept of concentration becomes ambiguous and should be replaced by activity. For a binary singly charged electrolyte KatAn = Kat + + An - the activities of the cation (a +) and anion (a -), respectively, are

a + = γ + . C + , a - = γ - . C - , (10.16)

where C + and C - are the analytical concentrations of the cation and anion, respectively;

γ + and γ - - their activity coefficients.

(10.17)

It is impossible to determine the activity of each ion separately, therefore, for singly charged electrolytes, the geometric mean values ​​of the activities i

and activity coefficients.

Instruction

The essence of this theory is that when melted (dissolved in water), almost all electrolytes decompose into ions, which are both positively and negatively charged (which is called electrolytic dissociation). Under the influence of an electric current, negative (“-”) towards the anode (+), and positively charged (cations, “+”), move towards the cathode (-). Electrolytic dissociation is a reversible process (the reverse process is called "molarization").

The degree (a) of electrolytic dissociation is dependent on the electrolyte itself, the solvent, and their concentration. This is the ratio of the number of molecules (n) that have decayed into ions to the total number of molecules introduced into the solution (N). You get: a = n / N

Thus, strong electrolytes are substances that completely decompose into ions when dissolved in water. Strong electrolytes, as a rule, are substances with highly polar or bonds: these are salts that are highly soluble (HCl, HI, HBr, HClO4, HNO3, H2SO4), as well as strong bases (KOH, NaOH, RbOH, Ba (OH) 2 , CsOH, Sr(OH)2, LiOH, Ca(OH)2). In a strong electrolyte, the substance dissolved in it is mostly in the form of ions ( ); there are practically no molecules that are undissociated.

Weak electrolytes are substances that only partially dissociate into ions. Weak electrolytes, along with ions in solution, contain undissociated molecules. Weak electrolytes do not give a strong concentration of ions in solution.

The weak ones are:
- organic acids (almost all) (C2H5COOH, CH3COOH, etc.);
- some of the acids (H2S, H2CO3, etc.);
- almost all salts, slightly soluble in water, ammonium hydroxide, as well as all bases (Ca3 (PO4) 2; Cu (OH) 2; Al (OH) 3; NH4OH);
- water.

They practically do not conduct electric current, or conduct, but poorly.

note

Although pure water conducts electricity very poorly, it still has a measurable electrical conductivity, due to the fact that water dissociates slightly into hydroxide ions and hydrogen ions.

Useful advice

Most electrolytes are corrosive substances, so when working with them, be extremely careful and follow safety regulations.

A strong base is an inorganic chemical compound formed by a hydroxyl group -OH and an alkali (group I elements of the periodic system: Li, K, Na, RB, Cs) or alkaline earth metal (group II elements Ba, Ca). They are written as formulas LiOH, KOH, NaOH, RbOH, CsOH, Ca(OH) ₂, Ba(OH) ₂.

You will need

• evaporating cup
• burner
• indicators
• metal rod
• H₃RO₄

Instruction

Strong bases exhibit, characteristic of all. The presence in the solution is determined by the change in color of the indicator. Add phenolphthalein to the sample with the test solution or omit litmus paper. Methyl orange is yellow, phenolphthalein is purple, and litmus paper is blue. The stronger the base, the more intense the color of the indicator.

If you need to find out which alkalis are presented to you, then conduct a qualitative analysis of the solutions. The most common strong bases are lithium, potassium, sodium, barium, and calcium. Bases react with acids (neutralization reactions) to form salt and water. In this case, Ca(OH) ₂, Ba(OH) ₂ and LiOH can be distinguished. When with acid, insoluble ones are formed. The remaining hydroxides will not give precipitation, tk. all K and Na salts are soluble.
3 Ca(OH) ₂ + 2 H₃RO₄ --→ Ca₃(PO₄)₂↓+ 6 H₂О

3 Va(OH) ₂ +2 H₃RO₄ --→ Va₃(PO₄)₂↓+ 6 H₂О

3 LiOH + Н₃РО₄ --→ Li₃РО₄↓ + 3 H₂О
Strain them and dry them. Inject the dried sediments into the flame of the burner. Lithium, calcium and barium ions can be qualitatively determined by changing the color of the flame. Accordingly, you will determine where which hydroxide is. Lithium salts color the burner flame carmine red. Barium salts - in green, and calcium salts - in raspberry.

The remaining alkalis form soluble orthophosphates.

3 NaOH + Н₃РО₄--→ Na₃РО₄ + 3 H₂О

3 KOH + H₃PO₄--→ K₃PO₄ + 3 H₂О

Evaporate the water to a dry residue. Evaporated salts on a metal rod alternately bring into the burner flame. There, sodium salt - the flame will turn bright yellow, and potassium - pink-purple. Thus, having a minimum set of equipment and reagents, you have determined all the strong reasons given to you.

An electrolyte is a substance that in the solid state is a dielectric, that is, does not conduct electric current, however, in a dissolved or molten form it becomes a conductor. Why is there such a drastic change in properties? The fact is that electrolyte molecules in solutions or melts dissociate into positively charged and negatively charged ions, due to which these substances in such a state of aggregation are able to conduct electric current. Most salts, acids, bases have electrolytic properties.

Instruction

What substances are strong? Such substances, in solutions or melts of which almost 100% of the molecules are exposed, and regardless of the concentration of the solution. The list includes the vast majority of soluble alkalis, salts and some acids, such as hydrochloric, bromine, iodine, nitric, etc.

And how do the weak ones behave in solutions or melts? electrolytes? Firstly, they dissociate to a very small extent (no more than 3% of the total number of molecules), and secondly, they go the worse and slower, the higher the concentration of the solution. Such electrolytes include, for example, (ammonium hydroxide), most organic and inorganic acids (including hydrofluoric - HF) and, of course, the familiar water to all of us. Since only a negligible fraction of its molecules decomposes into hydrogen ions and hydroxyl ions.

Remember that the degree of dissociation and, accordingly, the strength of the electrolyte are dependent on factors: the nature of the electrolyte itself, the solvent, and temperature. Therefore, this division itself is to a certain extent conditional. After all, the same substance can, under different conditions, be both a strong electrolyte and a weak one. To assess the strength of the electrolyte, a special value was introduced - the dissociation constant, determined on the basis of the law of mass action. But it is applicable only to weak electrolytes; strong electrolytes they do not obey the law of the acting masses.

Sources:

• strong electrolytes list

salt- These are chemicals consisting of a cation, that is, a positively charged ion, a metal and a negatively charged anion - an acid residue. There are many types of salts: normal, acidic, basic, double, mixed, hydrated, complex. It depends on the compositions of the cation and anion. How can you determine base salt?

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§ 6.3. Strong and weak electrolytes

The material in this section is partly familiar to you from previously studied school chemistry courses and from the previous section. Let's briefly review what you know and get acquainted with the new material.

In the previous section, we discussed the behavior in aqueous solutions of some salts and organic substances that completely decompose into ions in aqueous solution.
There is a number of simple but undoubted evidence that some substances in aqueous solutions decompose into particles. Thus, aqueous solutions of sulfuric H 2 SO 4 , nitric HNO 3 , chlorine HClO 4 , hydrochloric (hydrochloric) HCl, acetic CH 3 COOH and other acids have a sour taste. In the formulas of acids, the common particle is the hydrogen atom, and it can be assumed that (in the form of an ion) it is the cause of the same taste of all these so different substances.
The hydrogen ions formed during dissociation in an aqueous solution give the solution a sour taste, which is why such substances are called acids. In nature, only hydrogen ions taste sour. They create a so-called acidic (acidic) environment in an aqueous solution.

Remember, when you say “hydrogen chloride”, you mean the gaseous and crystalline state of this substance, but for an aqueous solution, you should say “hydrochloric acid solution”, “hydrochloric acid” or use the common name “hydrochloric acid”, although the composition of the substance in any state expressed by the same formula - Hcl.

Aqueous solutions of hydroxides of lithium (LiOH), sodium (NaOH), potassium (KOH), barium (Ba (OH) 2), calcium (Ca (OH) 2) and other metals have the same unpleasant bitter-soapy taste and cause on the skin of the hands sliding feeling. Apparently, OH– hydroxide ions, which are part of such compounds, are responsible for this property.
Hydrochloric HCl, hydrobromic HBr and hydroiodic HI acids react with zinc in the same way, despite their different composition, since it is not the acid that actually reacts with zinc:

Zn + 2HCl = ZnCl 2 + H2,

and hydrogen ions:

Zn + 2H + = Zn 2+ + H 2,

and hydrogen gas and zinc ions are formed.
The mixing of some salt solutions, for example, potassium chloride KCl and sodium nitrate NaNO 3, is not accompanied by a noticeable thermal effect, although after evaporation of the solution, a mixture of crystals of four substances is formed: the original ones - potassium chloride and sodium nitrate - and new ones - potassium nitrate KNO 3 and sodium chloride NaCl . It can be assumed that in a solution, the two initial salts completely decompose into ions, which, when it is evaporated, form four crystalline substances:

Comparing this information with the electrical conductivity of aqueous solutions of acids, hydroxides and salts and with a number of other provisions, S.A. Arrhenius in 1887 put forward the hypothesis of electrolytic dissociation, according to which the molecules of acids, hydroxides and salts, when dissolved in water, dissociate into ions.
The study of electrolysis products allows you to assign positive or negative charges to ions. Obviously, if an acid, for example, nitric HNO 3, dissociates, suppose, into two ions and hydrogen is released during the electrolysis of an aqueous solution on the cathode (negatively charged electrode), then, therefore, there are positively charged hydrogen ions H + in the solution. Then the dissociation equation should be written as follows:

HNO 3 \u003d H + +.

Electrolytic dissociation- complete or partial decomposition of the compound when it is dissolved in water into ions as a result of interaction with a water molecule (or other solvent).
electrolytes- acids, bases or salts, aqueous solutions of which conduct an electric current as a result of dissociation.
Substances that do not dissociate into ions in an aqueous solution and whose solutions do not conduct electricity are called non-electrolytes.
The dissociation of electrolytes is quantified degree of dissociation- the ratio of the number of "molecules" (formula units) decomposed into ions to the total number of "molecules" of the solute. The degree of dissociation is denoted by the Greek letter . For example, if out of every 100 "molecules" of a solute, 80 decompose into ions, then the degree of dissociation of the solute is: = 80/100 = 0.8, or 80%.
According to the ability to dissociate (or, as they say, “by strength”), electrolytes are divided into strong, medium and weak. According to the degree of dissociation, strong electrolytes include those for whose solutions > 30%, weak ones -< 3%, к средним – 3% 30%. Сила электролита – величина, зависящая от концентрации вещества, температуры, природы растворителя и др.
In the case of aqueous solutions, strong electrolytes(> 30%) belong to the following groups of compounds.
1 . Many inorganic acids, such as hydrochloric HCl, nitric HNO 3 , sulfuric H 2 SO 4 in dilute solutions. The strongest inorganic acid is perchloric HClO 4.
The strength of non-oxygen acids increases in a series of compounds of the same type when moving down the subgroup of acid-forming elements:

HCl-HBr-HI.

Hydrofluoric (hydrofluoric) acid HF dissolves glass, but this does not at all indicate its strength. This acid from oxygen-free halogen-containing acids belongs to acids of medium strength due to the high energy of the H–F bond, the ability of HF molecules to unite (associate) due to strong hydrogen bonds, the interaction of F ions with HF molecules (hydrogen bonds) with the formation of ions and other more complex particles. As a result, the concentration of hydrogen ions in an aqueous solution of this acid is significantly reduced, so hydrofluoric acid is considered to be of medium strength.
Hydrogen fluoride reacts with silicon dioxide, which is part of the glass, according to the equation:

SiO 2 + 4HF \u003d SiF 4 + 2H 2 O.

Hydrofluoric acid must not be stored in glass vessels. For this, vessels made of lead, some plastics and glass are used, the walls of which are covered from the inside with a thick layer of paraffin. If hydrogen fluoride gas is used to “etch” the glass, the glass surface becomes matte, which is used to apply inscriptions and various patterns on the glass. "Etching" the glass with an aqueous solution of hydrofluoric acid erodes the glass surface, which remains transparent. On sale is usually a 40% solution of hydrofluoric acid.

The strength of the same type of oxygen acids changes in the opposite direction, for example, iodic acid HIO 4 is weaker than perchloric acid HClO 4.
If an element forms several oxygen acids, then the acid in which the acid-forming element has the highest valence has the greatest strength. So, in the series of acids HclO (hypochlorous) - HclO 2 (chloric) - HclO 3 (chloric) - HclO 4 (chloric), the latter is the strongest.

One volume of water dissolves about two volumes of chlorine. Chlorine (about half of it) interacts with water:

Cl 2 + H 2 O \u003d HCl + HClO.

Hydrochloric acid is strong; there are practically no HCl molecules in its aqueous solution. The correct equation for the reaction is:

Cl 2 + H 2 O \u003d H + + Cl - + HClO - 25 kJ / mol.

The resulting solution is called chlorine water.
Hypochlorous acid is a fast-acting oxidizing agent, so it is used to bleach fabrics.

2 . Hydroxides of elements of the main subgroups of groups I and II of the periodic system: LiOH, NaOH, KOH, Ca (OH) 2, etc. When moving down the subgroup, as the metallic properties of the element increase, the strength of the hydroxides increases. Soluble hydroxides of the main subgroup of group I elements are classified as alkalis.

Bases soluble in water are called alkalis. These also include the hydroxides of the elements of the main subgroup of group II (alkaline earth metals) and ammonium hydroxide (an aqueous solution of ammonia). Sometimes alkalis are those hydroxides that create a high concentration of hydroxide ions in an aqueous solution. In outdated literature, you can find among the alkalis potassium carbonates K 2 CO 3 (potash) and sodium Na 2 CO 3 (soda), sodium bicarbonate NaHCO 3 (baking soda), borax Na 2 B 4 O 7, sodium hydrosulfides NaHS and potassium KHS etc.

Calcium hydroxide Ca (OH) 2 as a strong electrolyte dissociates into one step:

Ca (OH) 2 \u003d Ca 2+ + 2OH -.

3 . Almost all salts. Salt, if it is a strong electrolyte, dissociates into one step, for example ferric chloride:

FeCl 3 \u003d Fe 3+ + 3Cl -.

In the case of aqueous solutions, weak electrolytes ( < 3%) относят перечисленные ниже соединения.

1 . Water H 2 O is the most important electrolyte.

2 . Some inorganic and almost all organic acids: H 2 S (hydrosulfide), H 2 SO 3 (sulphurous), H 2 CO 3 (carbonic), HCN (hydrocyanic), H 3 PO 4 (phosphoric, orthophosphoric), H 2 SiO 3 (silicon), H 3 BO 3 (boric, orthoboric), CH 3 COOH (acetic), etc.
Note that carbonic acid does not exist in the formula H 2 CO 3. When carbon dioxide CO 2 is dissolved in water, its hydrate CO 2 H 2 O is formed, which we write for convenience of calculations by the formula H 2 CO 3, and the equation for the dissociation reaction looks like this:

The dissociation of weak carbonic acid proceeds in two steps. The resulting bicarbonate ion also behaves like a weak electrolyte.
Other polybasic acids dissociate in the same way: H 3 PO 4 (phosphoric), H 2 SiO 3 (silicon), H 3 BO 3 (boric). In an aqueous solution, dissociation practically passes only through the first stage. How to carry out dissociation along the last step?
3 . Hydroxides of many elements, such as Al (OH) 3, Cu (OH) 2, Fe (OH) 2, Fe (OH) 3, etc.
All these hydroxides dissociate in an aqueous solution in steps, for example, iron hydroxide
Fe(OH)3:

In an aqueous solution, dissociation proceeds practically only through the first stage. How to shift the equilibrium towards the formation of Fe 3+ ions?
The basic properties of hydroxides of the same element increase with a decrease in the valency of the element. Thus, the basic properties of iron dihydroxide Fe (OH) 2 are more pronounced than those of Fe (OH) 3 trihydroxide. This statement is equivalent to the fact that the acidic properties of Fe(OH) 3 are stronger than that of Fe(OH) 2 .
4 . Ammonium hydroxide NH 4 OH.
When gaseous ammonia NH 3 is dissolved in water, a solution is obtained that conducts electricity very poorly and has a bitter-soapy taste. The solution medium is basic or alkaline. This behavior of ammonia is explained as follows. When ammonia is dissolved in water, ammonia hydrate NH 3 H 2 O is formed, to which we conditionally attribute the formula of the non-existent ammonium hydroxide NH 4 OH, assuming that this compound dissociates with the formation ammonium ion and hydroxide ion OH -:

NH 4 OH \u003d + OH -.

5 . Some salts: zinc chloride ZnCl 2, iron thiocyanate Fe (NCS) 3, mercury cyanide Hg (CN) 2, etc. These salts dissociate in steps.

For electrolytes of medium strength, some include phosphoric acid H 3 PO 4. We will consider phosphoric acid as a weak electrolyte and write down the three steps of its dissociation. Sulfuric acid in concentrated solutions behaves like an electrolyte of medium strength, and in very concentrated solutions it behaves like a weak electrolyte. We will further consider sulfuric acid as a strong electrolyte and write the equation for its dissociation in one step.

Strong electrolytes, when dissolved in water, almost completely dissociate into ions, regardless of their concentration in solution.

Therefore, in the equations of dissociation of strong electrolytes put an equal sign (=).

Strong electrolytes include:

Soluble salts;

Many inorganic acids: HNO3, H2SO4, HCl, HBr, HI;

Bases formed by alkali metals (LiOH, NaOH, KOH, etc.) and alkaline earth metals (Ca(OH)2, Sr(OH)2, Ba(OH)2).

Weak electrolytes in aqueous solutions only partially (reversibly) dissociate into ions.

Therefore, the reversibility sign (⇄) is put in the dissociation equations for weak electrolytes.

Weak electrolytes include:

Almost all organic acids and water;

Some inorganic acids: H2S, H3PO4, H2CO3, HNO2, H2SiO3, etc.;

Insoluble metal hydroxides: Mg(OH)2, Fe(OH)2, Zn(OH)2, etc.

Ionic reaction equations

Ionic reaction equations
Chemical reactions in electrolyte solutions (acids, bases and salts) proceed with the participation of ions. The final solution may remain transparent (the products are highly soluble in water), but one of the products will turn out to be a weak electrolyte; in other cases, precipitation or gas evolution will be observed.

For reactions in solutions involving ions, not only the molecular equation is compiled, but also the full ionic and short ionic equations.
In ionic equations, at the suggestion of the French chemist K.-L. Berthollet (1801), all strong, well-soluble electrolytes are written in the form of ion formulas, and precipitation, gases and weak electrolytes are written in the form of molecular formulas. The formation of precipitation is marked with a down arrow sign (↓), the formation of gases with an up arrow sign (). An example of writing the reaction equation according to the Berthollet rule:

a) molecular equation
Na2CO3 + H2SO4 = Na2SO4 + CO2 + H2O
b) complete ionic equation
2Na+ + CO32− + 2H+ + SO42− = 2Na+ + SO42− + CO2 + H2O
(CO2 - gas, H2O - weak electrolyte)
c) short ionic equation
CO32− + 2H+ = CO2 + H2O

Usually, when writing, they are limited to a brief ionic equation, with solid reagents denoted by the index (t), gaseous reagents - by the index (g). Examples:

1) Cu(OH)2(t) + 2HNO3 = Cu(NO3)2 + 2H2O
Cu(OH)2(t) + 2H+ = Cu2+ + 2H2O
Cu(OH)2 is practically insoluble in water
2) BaS + H2SO4 = BaSO4↓ + H2S
Ba2+ + S2− + 2H+ + SO42− = BaSO4↓ + H2S
(full and short ionic equations are the same)
3) CaCO3(t) + CO2(g) + H2O = Ca(HCO3)2
CaCO3(t) + CO2(g) + H2O = Ca2+ + 2HCO3−
(most acid salts are highly soluble in water).

If strong electrolytes do not participate in the reaction, there is no ionic form of the equation:

Mg(OH)2(t) + 2HF(p) = MgF2↓ + 2H2O

TICKET #23

Salt hydrolysis

Salt hydrolysis is the interaction of salt ions with water to form low-dissociating particles.

Hydrolysis, literally, is the decomposition by water. By giving this definition of the reaction of hydrolysis of salts, we emphasize that salts in solution are in the form of ions, and that the driving force of the reaction is the formation of slightly dissociating particles (a general rule for many reactions in solutions).

Hydrolysis occurs only in those cases when the ions formed as a result of the electrolytic dissociation of the salt - a cation, an anion, or both together - are able to form weakly dissociating compounds with water ions, and this, in turn, occurs when the cation is strongly polarizing ( weak base cation), and the anion is easily polarized (weak acid anion). This changes the pH of the medium. If the cation forms a strong base, and the anion forms a strong acid, then they do not undergo hydrolysis.

1. Hydrolysis of a salt of a weak base and a strong acid passes through the cation, this may form a weak base or basic salt and the pH of the solution will decrease

2. Hydrolysis of a salt of a weak acid and a strong base passes through the anion, a weak acid or an acidic salt may be formed and the pH of the solution will increase

3. Hydrolysis of a salt of a weak base and a weak acid usually passes through to form a weak acid and a weak base; The pH of the solution in this case differs slightly from 7 and is determined by the relative strength of the acid and base

4. Hydrolysis of a salt of a strong base and a strong acid does not proceed

Question 24 Classification of oxides

Oxides complex substances are called, the composition of the molecules of which includes oxygen atoms in the oxidation state - 2 and some other element.

oxides can be obtained by direct interaction of oxygen with another element, or indirectly (for example, by the decomposition of salts, bases, acids). Under normal conditions, oxides are in a solid, liquid and gaseous state, this type of compounds is very common in nature. Oxides are found in the Earth's crust. Rust, sand, water, carbon dioxide are oxides.

Salt-forming oxides For example,

CuO + 2HCl → CuCl 2 + H 2 O.

CuO + SO 3 → CuSO 4.

Salt-forming oxides- These are oxides that form salts as a result of chemical reactions. These are oxides of metals and non-metals, which, when interacting with water, form the corresponding acids, and when interacting with bases, the corresponding acidic and normal salts. For example, copper oxide (CuO) is a salt-forming oxide, because, for example, when it reacts with hydrochloric acid (HCl), a salt is formed:

CuO + 2HCl → CuCl 2 + H 2 O.

As a result of chemical reactions, other salts can be obtained:

CuO + SO 3 → CuSO 4.

Non-salt-forming oxides called oxides that do not form salts. An example is CO, N 2 O, NO.

ELECTROLYTES Substances whose solutions or melts conduct electricity.

NON-ELECTROLYTES Substances whose solutions or melts do not conduct electricity.

Dissociation- decomposition of compounds into ions.

Degree of dissociation is the ratio of the number of molecules dissociated into ions to the total number of molecules in the solution.

STRONG ELECTROLYTES when dissolved in water, they almost completely dissociate into ions.

When writing the equations of dissociation of strong electrolytes put an equal sign.

Strong electrolytes include:

Soluble salts ( see solubility table);

Many inorganic acids: HNO 3, H 2 SO 4, HClO 3, HClO 4, HMnO 4, HCl, HBr, HI ( look acids-strong electrolytes in the solubility table);

Bases of alkali (LiOH, NaOH, KOH) and alkaline earth (Ca (OH) 2, Sr (OH) 2, Ba (OH) 2) metals ( see strong electrolyte bases in the solubility table).

WEAK ELECTROLYTES in aqueous solutions only partially (reversibly) dissociate into ions.

When writing the dissociation equations for weak electrolytes, the sign of reversibility is put.

Weak electrolytes include:

Almost all organic acids and water (H 2 O);

Some inorganic acids: H 2 S, H 3 PO 4, HClO 4, H 2 CO 3, HNO 2, H 2 SiO 3 ( look acids-weak electrolytes in the solubility table);

Insoluble metal hydroxides (Mg (OH) 2, Fe (OH) 2, Zn (OH) 2) ( see basescweak electrolytes in the solubility table).

The degree of electrolytic dissociation is influenced by a number of factors:

the nature of the solvent and electrolyte: strong electrolytes are substances with ionic and covalent strongly polar bonds; good ionizing ability, i.e. the ability to cause dissociation of substances, have solvents with a high dielectric constant, the molecules of which are polar (for example, water);

temperature: since dissociation is an endothermic process, an increase in temperature increases the value of α;

concentration: when the solution is diluted, the degree of dissociation increases, and with increasing concentration, it decreases;

stage of the dissociation process: each subsequent stage is less effective than the previous one, approximately 1000–10,000 times; for example, for phosphoric acid α 1 > α 2 > α 3:

H3PO4⇄Н++H2PO−4 (first stage, α 1),

H2PO−4⇄H++HPO2−4 (second stage, α 2),

НPO2−4⇄Н++PO3−4 (third stage, α 3).

For this reason, in a solution of this acid, the concentration of hydrogen ions is the highest, and the concentration of PO3−4 phosphate ions is the lowest.

1. Solubility and the degree of dissociation of a substance are not related to each other. For example, a weak electrolyte is acetic acid, which is highly (unrestrictedly) soluble in water.

2. A solution of a weak electrolyte contains less than others those ions that are formed at the last stage of electrolytic dissociation

The degree of electrolytic dissociation is also affected by addition of other electrolytes: e.g. degree of dissociation of formic acid

HCOOH ⇄ HCOO − + H+

decreases if a little sodium formate is added to the solution. This salt dissociates to form formate ions HCOO − :

HCOONa → HCOO − + Na +

As a result, the concentration of HCOO– ions in the solution increases, and according to the Le Chatelier principle, an increase in the concentration of formate ions shifts the equilibrium of the formic acid dissociation process to the left, i.e. the degree of dissociation decreases.

Ostwald dilution law- ratio expressing the dependence of the equivalent electrical conductivity of a dilute solution of a binary weak electrolyte on the concentration of the solution:

Here, is the dissociation constant of the electrolyte, is the concentration, and are the values ​​of the equivalent electrical conductivity at concentration and at infinite dilution, respectively. The ratio is a consequence of the law of mass action and equality

where is the degree of dissociation.

The Ostwald dilution law was developed by W. Ostwald in 1888 and confirmed by him experimentally. The experimental establishment of the correctness of the Ostwald dilution law was of great importance for substantiating the theory of electrolytic dissociation.

Electrolytic dissociation of water. Hydrogen indicator pH Water is a weak amphoteric electrolyte: H2O H+ + OH- or, more precisely: 2H2O \u003d H3O + + OH- The dissociation constant of water at 25 ° C is: can be considered constant and equal to 55.55 mol / l (water density 1000 g / l, mass 1 l 1000 g, amount of water substance 1000g: 18g / mol \u003d 55.55 mol, C \u003d 55.55 mol: 1 l \u003d 55 .55 mol/l). Then This value is constant at a given temperature (25 ° C), it is called the ion product of water KW: The dissociation of water is an endothermic process, therefore, with an increase in temperature, in accordance with the Le Chatelier principle, dissociation increases, the ion product increases and reaches a value of 10-13 at 100 ° C. In pure water at 25°C, the concentrations of hydrogen and hydroxyl ions are equal to each other: = = 10-7 mol/l Solutions in which the concentrations of hydrogen and hydroxyl ions are equal to each other are called neutral. If acid is added to pure water, the concentration of hydrogen ions will increase and become more than 10-7 mol / l, the medium will become acidic, while the concentration of hydroxyl ions will instantly change so that the ion product of water retains its value of 10-14. The same thing will happen when alkali is added to pure water. The concentrations of hydrogen and hydroxyl ions are related to each other through the ion product, therefore, knowing the concentration of one of the ions, it is easy to calculate the concentration of the other. For example, if = 10-3 mol/l, then = KW/ = 10-14/10-3 = 10-11 mol/l, or if = 10-2 mol/l, then = KW/ = 10-14 /10-2 = 10-12 mol/l. Thus, the concentration of hydrogen or hydroxyl ions can serve as a quantitative characteristic of the acidity or alkalinity of the medium. In practice, it is not the concentrations of hydrogen or hydroxyl ions that are used, but the hydrogen pH or hydroxyl pOH indicators. The hydrogen index pH is equal to the negative decimal logarithm of the concentration of hydrogen ions: pH = - lg The hydroxyl index pOH is equal to the negative decimal logarithm of the concentration of hydroxyl ions: pOH = - lg It is easy to show by pronouncing the ionic product of water that pH + pOH = 14 the medium is neutral, if less than 7 - acidic, and the lower the pH, the higher the concentration of hydrogen ions. pH greater than 7 - alkaline environment, the higher the pH, the higher the concentration of hydroxyl ions.