Weak electrolytes- substances that partially dissociate into ions. Solutions of weak electrolytes contain undissociated molecules along with ions. Weak electrolytes cannot produce 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 salts, bases and ammonium hydroxide Ca 3 (PO 4) 2 that are slightly soluble in water; Cu(OH) 2 ; Al(OH) 3 ; NH4OH;

They conduct electricity poorly (or almost not at all).

The concentrations of ions 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 = 0.3 is the conventional boundary for dividing 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) per solute molecule increases. According to Le Chatelier’s principle, the equilibrium of electrolytic dissociation in this case should shift in the direction of the formation of products, i.e. hydrated ions.

The degree of electrolytic dissociation depends on the temperature of the solution. Typically, as the temperature increases, the degree of dissociation increases, because bonds in molecules are activated, they become more mobile and are easier to ionize. The concentration of ions in a weak electrolyte solution can be calculated by knowing the degree of dissociation a and 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 the weak electrolyte C and its degree of dissociation α, we obtain:

K d = C. α. S. α/S. (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 = C. α 2. (10.13)

This allows us to conclude that with 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 measurement and recording, the value used is the hydrogen index, (criterion) of the strength of an acid or base. By analogy .

From equation (11.15): . At pH=7 – the solution reaction 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 Strong electrolyte solutions

Strong electrolytes are substances that, when dissolved in water, almost completely disintegrate 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 strong electrolyte solution, the solute is found primarily in the form of ions (cations and anions); undissociated molecules are practically absent.

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

H 2 SO 4 = H + + HSO 4 - ,

and therefore the equilibrium (dissociation) constant turns out to be an uncertain 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, having proposed a model that formed the basis of the theory of strong electrolytes, 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 an ionic atmosphere. During 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 in front of him and thins out behind him. This asymmetry of the ionic atmosphere has a more inhibiting 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 must be replaced by activity. For a binary single-charge electrolyte KatAn = Kat + + An - the activities of the cation (a +) and anion (a -) are respectively equal

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 - are their activity coefficients.

(10.17)

It is impossible to determine the activity of each ion separately; therefore, for single-charge electrolytes, geometric mean values ​​of the activities are used.

and activity coefficients:

The Debye-Hückel activity coefficient depends at least on temperature, dielectric constant of the solvent (ε), and ionic strength (I); the latter serves as a measure of the intensity of the electric field created by the ions in the solution.

For a given electrolyte, ionic strength is expressed by the Debye-Hückel equation:

The ionic strength in turn is equal to

where C is the analytical concentration;

z is the charge of the cation or anion.

For a singly charged electrolyte, the ionic strength coincides with the concentration. Thus, NaCl and Na 2 SO 4 at the same concentrations will have different ionic strengths. Comparison of the properties of solutions of strong electrolytes can only be carried out when the ionic strengths are the same; even small impurities dramatically change the properties of the electrolyte.

Figure 10.5 - Dependency

Electrolyte dissociation is quantitatively characterized by the degree of dissociation. Dissociation degree a–this is the ratio of the number of molecules dissociated into ions N diss.,to the total number of molecules of dissolved electrolyte N :

a =

a– the fraction of electrolyte molecules that have broken up into ions.

The degree of electrolyte dissociation depends on many factors: the nature of the electrolyte, the nature of the solvent, the concentration of the solution, and temperature.

Based on their ability to dissociate, electrolytes are conventionally divided into strong and weak. Electrolytes that exist in solution only in the form of ions are usually called strong . Electrolytes, which in a dissolved state are partly in the form of molecules and partly in the form of ions, are called weak .

Strong electrolytes include almost all salts, some acids: H 2 SO 4, HNO 3, HCl, HI, HClO 4, hydroxides of alkali and alkaline earth metals (see appendix, table 6).

The process of dissociation of strong electrolytes continues to completion:

HNO 3 = H + + NO 3 - , NaOH = Na + + OH - ,

and equal signs are placed in the dissociation equations.

In relation to strong electrolytes, the concept of “degree of dissociation” is conditional. " Apparent degree of dissociation (a each) below the true one (see appendix, table 6). With increasing concentration of a strong electrolyte in a solution, the interaction of oppositely charged ions increases. When sufficiently close to each other, they form associates. The ions in them are separated by layers of polar water molecules surrounding each ion. This affects the decrease in the electrical conductivity of the solution, i.e. the effect of incomplete dissociation is created.

To take this effect into account, an activity coefficient g was introduced, which decreases with increasing concentration of the solution, varying from 0 to 1. To quantitatively describe the properties of solutions of strong electrolytes, a quantity called activity (a).

The activity of an ion is understood as its effective concentration, according to which it acts in chemical reactions.

Ion activity ( a) is equal to its molar concentration ( WITH), multiplied by the activity coefficient (g):

A = g WITH.

Using activity instead of concentration allows one to apply to solutions the laws established for ideal solutions.

Weak electrolytes include some mineral acids (HNO 2, H 2 SO 3, H 2 S, H 2 SiO 3, HCN, H 3 PO 4) and most organic acids (CH 3 COOH, H 2 C 2 O 4, etc.) , ammonium hydroxide NH 4 OH and all bases that are slightly soluble in water, organic amines.

The dissociation of weak electrolytes is reversible. In solutions of weak electrolytes, an equilibrium is established between ions and undissociated molecules. In the corresponding dissociation equations, the reversibility sign (“”) is placed. For example, the dissociation equation for weak acetic acid is written as follows:

CH 3 COOH « CH 3 COO - + H + .

In a solution of a weak binary electrolyte ( CA) the following equilibrium is established, characterized by an equilibrium constant called the dissociation constant TO d:

KA « K + + A - ,

.

If 1 liter of solution is dissolved WITH moles of electrolyte CA and the degree of dissociation is a, which means dissociated aС moles of electrolyte and each ion was formed aС moles. In the undissociated state remains ( WITH – aС) moles CA.

KA « K + + A - .

C – aС aС aС

Then the dissociation constant will be equal to:

(6.1)

Since the dissociation constant does not depend on concentration, the derived relation expresses the dependence of the degree of dissociation of a weak binary electrolyte on its concentration. From equation (6.1) it is clear that a decrease in the concentration of a weak electrolyte in a solution leads to an increase in the degree of its dissociation. Equation (6.1) expresses Ostwald's dilution law .

For very weak electrolytes (at a<<1), уравнение Оствальда можно записать следующим образом:

TO d a 2 C, or a" (6.2)

The dissociation constant for each electrolyte is constant at a given temperature, it does not depend on the concentration of the solution and characterizes the ability of the electrolyte to disintegrate into ions. The higher the Kd, the more the electrolyte dissociates into ions. The dissociation constants of weak electrolytes are tabulated (see appendix, table 3).

Instructions

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

The degree of (a) electrolytic dissociation depends on the electrolyte itself, the solvent, and their concentration. This is the ratio of the number of molecules (n) that broke up 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 disintegrate into ions when dissolved in water. Strong electrolytes are usually 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 dissociate into ions only partially. Weak electrolytes, together with ions in solution, contain undissociated molecules. Weak electrolytes do not produce a strong concentration of ions in solution.

The weak ones include:
- organic acids (almost all) (C2H5COOH, CH3COOH, etc.);
- some of the acids (H2S, H2CO3, etc.);
- almost all salts that are 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.

Please note

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

Useful advice

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

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

You will need

• evaporation cup
• burner
• indicators
• metal rod
• N₃PO₄

Instructions

Strong reasons are manifested, 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 the litmus paper. Methyl orange produces a yellow color, phenolphthalein produces a purple color, and litmus paper turns 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 combined with acid, insoluble compounds are formed. The remaining hydroxides will not produce precipitation, because All K and Na salts are soluble.
3 Ca(OH) ₂ + 2 H₃PO₄ --→ Ca₃(PO₄)₂↓+ 6 H₂O

3 Ba(OH) ₂ +2 Н₃PO₄ --→ Ba₃(PO₄)₂↓+ 6 H₂О

3 LiOH + H₃PO₄ --→ Li₃PO₄↓ + 3 H₂O
Strain them and dry them. Add the dried sediment to the burner flame. By changing the color of the flame, lithium, calcium and barium ions can be qualitatively determined. Accordingly, you will determine which hydroxide is which. Lithium salts color the burner flame carmine red. Barium salts are green, and calcium salts are crimson.

The remaining alkalis form soluble orthophosphates.

3 NaOH + H₃PO₄--→ Na₃PO₄ + 3 H₂O

3 KOH + H₃PO₄--→ K₃PO₄ + 3 H₂O

It is necessary to evaporate the water to a dry residue. Place the evaporated salts on a metal rod one by one into the burner flame. There, sodium salt - the flame will turn bright yellow, and potassium - pink-violet. Thus, having a minimal set of equipment and reagents, you have identified all the strong reasons given to you.

An electrolyte is a substance that in its solid state is a dielectric, that is, it does not conduct electric current, but when dissolved or molten it becomes a conductor. Why does such a sharp change in properties occur? 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 an aggregate state are capable of conducting electric current. Most salts, acids, and bases have electrolytic properties.

Instructions

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

How do 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, their dissociation becomes worse and slower the higher the concentration of the solution. Such electrolytes include, for example, (ammonium hydroxide), most organic and inorganic acids (including hydrofluoric acid - HF) and, of course, familiar water to all of us. Since only a negligible fraction of its molecules breaks down into hydrogen ions and hydroxyl ions.

Remember that the degree of dissociation and, accordingly, the strength of the electrolyte depend on factors: the nature of the electrolyte itself, the solvent, and temperature. Therefore, this division itself is to a certain extent arbitrary. 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 do not obey the law of mass action.

Sources:

• strong electrolytes list

Salts- these are chemical substances 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. This depends on the cation and anion compositions. How can you determine base salt?

1. ELECTROLYTES

1.1. Electrolytic dissociation. Degree of dissociation. Electrolyte Power

According to the theory of electrolytic dissociation, salts, acids, and hydroxides, when dissolved in water, completely or partially disintegrate into independent particles - ions.

The process of decomposition of substance molecules into ions under the influence of polar solvent molecules is called electrolytic dissociation. Substances that dissociate into ions in solutions are called electrolytes. As a result, the solution acquires the ability to conduct electric current, because mobile electric charge carriers appear in it. According to this theory, when dissolved in water, electrolytes break up (dissociate) into positively and negatively charged ions. Positively charged ions are called cations; these include, for example, hydrogen and metal ions. Negatively charged ions are called anions; These include ions of acidic residues and hydroxide ions.

To quantitatively characterize the dissociation process, the concept of the degree of dissociation was introduced. The degree of dissociation of an electrolyte (α) is the ratio of the number of its molecules disintegrated into ions in a given solution ( n ), to the total number of its molecules in solution ( N), or

α = .

The degree of electrolytic dissociation is usually expressed either in fractions of a unit or as a percentage.

Electrolytes with a degree of dissociation greater than 0.3 (30%) are usually called strong, with a degree of dissociation from 0.03 (3%) to 0.3 (30%) - medium, less than 0.03 (3%) - weak electrolytes. So, for a 0.1 M solution CH3COOH α = 0.013 (or 1.3%). Therefore, acetic acid is a weak electrolyte. The degree of dissociation shows what part of the dissolved molecules of a substance has broken up into ions. The degree of electrolytic dissociation of an electrolyte in aqueous solutions depends on the nature of the electrolyte, its concentration and temperature.

By their nature, electrolytes can be divided into two large groups: strong and weak. Strong electrolytes dissociate almost completely (α = 1).

Strong electrolytes include:

1) acids (H 2 SO 4, HCl, HNO 3, HBr, HI, HClO 4, H M nO 4);

2) bases – metal hydroxides of the first group of the main subgroup (alkali) – LiOH, NaOH, KOH, RbOH, CsOH , as well as hydroxides of alkaline earth metals – Ba (OH) 2, Ca (OH) 2, Sr (OH) 2;.

3) salts soluble in water (see solubility table).

Weak electrolytes dissociate into ions to a very small extent; in solutions they are found mainly in an undissociated state (in molecular form). For weak electrolytes, an equilibrium is established between undissociated molecules and ions.

Weak electrolytes include:

1) inorganic acids ( H 2 CO 3, H 2 S, HNO 2, H 2 SO 3, HCN, H 3 PO 4, H 2 SiO 3, HCNS, HClO, etc.);

2) water (H 2 O);

3) ammonium hydroxide ( NH 4 OH);

4) most organic acids

(for example, acetic CH 3 COOH, formic HCOOH);

5) insoluble and slightly soluble salts and hydroxides of some metals (see solubility table).

Process electrolytic dissociation depicted using chemical equations. For example, dissociation of hydrochloric acid (HC l ) is written as follows:

HCl → H + + Cl – .

Bases dissociate to form metal cations and hydroxide ions. For example, the dissociation of KOH

KOH → K + + OH – .

Polybasic acids, as well as bases of polyvalent metals, dissociate stepwise. For example,

H 2 CO 3 H + + HCO 3 – ,

HCO 3 – H + + CO 3 2– .

The first equilibrium - dissociation according to the first step - is characterized by the constant

.

For second stage dissociation:

.

In the case of carbonic acid, the dissociation constants have the following values: K I = 4.3× 10 –7, K II = 5.6 × 10–11. For stepwise dissociation always K I > K II > K III >... , because the energy that must be expended to separate an ion is minimal when it is separated from a neutral molecule.

Average (normal) salts, soluble in water, dissociate to form positively charged metal ions and negatively charged ions of the acid residue

Ca(NO 3) 2 → Ca 2+ + 2NO 3 –

Al 2 (SO 4) 3 → 2Al 3+ +3SO 4 2–.

Acid salts (hydrosalts) are electrolytes containing hydrogen in the anion, which can be split off in the form of the hydrogen ion H +. Acid salts are considered as a product obtained from polybasic acids in which not all hydrogen atoms are replaced by a metal. Dissociation of acid salts occurs in stages, for example:

KHCO 3 → K + + HCO 3 – (first stage)

There are strong and weak electrolytes. Strong electrolytes in solutions are almost completely dissociated. This group of electrolytes includes most salts, alkalis and strong acids. Weak electrolytes include weak acids and weak bases and some salts: mercury (II) chloride, mercury (II) cyanide, iron (III) thiocyanate, cadmium iodide. Solutions of strong electrolytes at high concentrations have significant electrical conductivity, and it increases slightly with dilution of solutions.

Solutions of weak electrolytes at high concentrations are characterized by insignificant electrical conductivity, which increases greatly when the solutions are diluted.

When a substance is dissolved in any solvent, simple (unsolvated) ions, neutral molecules of the dissolved substance, solvated (hydrated in aqueous solutions) ions (for example, etc.), ion pairs (or ion twins), which are electrostatically associated groups of oppositely charged ions (for example, ), the formation of which is observed in the overwhelming majority of non-aqueous electrolyte solutions, complex ions (for example, ), solvated molecules, etc.

In aqueous solutions of strong electrolytes, only simple or solvated cations and anions exist. There are no solute molecules in their solutions. Therefore, it is incorrect to assume the presence of molecules or the presence of long-term bonds between or and in an aqueous solution of sodium chloride.

In aqueous solutions of weak electrolytes, the solute can exist in the form of simple and solvated (-hydrated) ions and undissociated molecules.

In non-aqueous solutions, some strong electrolytes (for example, ) are not completely dissociated even at moderately high concentrations. In most organic solvents, the formation of ion pairs of oppositely charged ions is observed (for more details, see book 2).

In some cases, it is impossible to draw a sharp line between strong and weak electrolytes.

Interional forces. Under the influence of interionic forces, around each freely moving ion, other ions charged with the opposite sign are grouped, arranged symmetrically, forming the so-called ionic atmosphere, or ion cloud, slowing down the movement of the ion in the solution.

For example, in a solution, chlorine ions are grouped around moving potassium ions, and an atmosphere of potassium ions is created near the moving chlorine ions.

Ions whose mobility is weakened by interionic extension forces exhibit reduced chemical activity in solutions. This causes deviations in the behavior of strong electrolytes from the classical form of the law of mass action.

Foreign ions present in a given electrolyte solution also have a strong effect on the mobility of its ions. The higher the concentration, the more significant the interionic interaction and the more strongly the foreign ions affect the mobility of the ions.

In weak acids and bases, the hydrogen or hydroxyl bond in their molecules is largely covalent rather than ionic; Therefore, when weak electrolytes are dissolved in solvents characterized by a very high dielectric constant, most of their molecules do not disintegrate into ions.

Solutions of strong electrolytes differ from solutions of weak electrolytes in that they do not contain undissociated molecules. This is confirmed by modern physical and physicochemical studies. For example, X-ray examination of crystals of strong electrolytes confirms the fact that the crystal lattices of salts are built from ions.

When dissolved in a solvent with a high dielectric constant, solvate shells (hydrate in water) are formed around the ions, preventing them from combining into molecules. Thus, since strong electrolytes do not contain molecules even in the crystalline state, they especially do not contain molecules in solutions.

However, it was found experimentally that the electrical conductivity of aqueous solutions of strong electrolytes is not equivalent to the electrical conductivity that could be expected during the dissociation of dissolved electrolyte molecules into ions.

Using the theory of electrolytic dissociation proposed by Arrhenius, it turned out to be impossible to explain this and a number of other facts. To explain them, new scientific principles were put forward.

At present, the discrepancy between the properties of strong electrolytes and the classical form of the law of mass action can be explained using the theory of strong electrolytes proposed by Debye and Hückel. The main idea of ​​this theory is that mutual attractive forces arise between ions of strong electrolytes in solutions. These interionic forces cause the behavior of strong electrolytes to deviate from the laws of ideal solutions. The presence of these interactions causes mutual inhibition of cations and anions.

Effect of dilution on interionic attraction. Interionic attraction causes deviations in the behavior of real solutions in the same way as intermolecular attraction in real gases entails deviations in their behavior from the laws of ideal gases. The higher the concentration of the solution, the denser the ionic atmosphere and the lower the mobility of ions, and therefore the electrical conductivity of electrolytes.

Just as the properties of a real gas at low pressures approach the properties of an ideal gas, so the properties of solutions of strong electrolytes at high dilutions approach the properties of ideal solutions.

In other words, in dilute solutions the distances between the ions are so large that the mutual attraction or repulsion experienced by the ions is extremely small and practically reduced to zero.

Thus, the observed increase in the electrical conductivity of strong electrolytes when their solutions are diluted is explained by the weakening of interionic forces of attraction and repulsion, which causes an increase in the speed of movement of ions.

The less dissociated the electrolyte and the more dilute the solution, the smaller the interionic electrical influence and the fewer deviations from the law of mass action are observed, and, conversely, the higher the concentration of the solution, the greater the interionic electrical influence and the more deviations from the law of mass action are observed.

For the reasons stated above, the law of mass action in its classical form cannot be applied to aqueous solutions of strong electrolytes, as well as to concentrated aqueous solutions of weak electrolytes.