35. Degree and dissociation constant of a weak electrolyte. Ostwald's law of dilution. Stepwise dissociation of the electrolyte. The influence of common ions on the dissociation of weak electrolytes.
Degree of dissociation (alpha) of an electrolyte is the proportion of its molecules that undergo dissociation.
Dissociation constant- a type of equilibrium constant that shows the tendency of a large object to dissociate (separate) reversibly into small objects, such as when a complex breaks down into its component molecules, or when a salt separates into ions in an aqueous solution.
Oswald's law of dilution:K= Cm/(1-α)
Polybasic acids, as well as bases of two or more valent metals dissociate stepwise. In solutions of these substances, complex equilibria are established in which ions of different charges participate.
First equilibrium - first stage dissociation– characterized by a dissociation constant, denoted TO 1 , and the second - second stage dissociation – dissociation constant TO 2 . Quantities K, K 1 And TO 2 are related to each other by the relation: K=K 1 TO 2
During the stepwise dissociation of substances, the decomposition in the subsequent step always occurs to a lesser extent than in the previous one. The following inequality holds: TO 1 >K 2 >K 3 …
This is explained by the fact that the energy that must be expended to remove an ion is minimal when it is separated from a neutral molecule and becomes greater during dissociation at each subsequent step.
Effect of a common ion on the dissociation of a weak electrolyte: Addition of a common ion reduces the dissociation of the weak electrolyte.
36. Self-ionization of water. Ionic product of water. Hydrogen (pH) and hydroxyl (pOH) indicators, their relationship in water and aqueous solutions of electrolytes. The concept of indicators and buffer solutions of electrolytes. Concept of indicators and buffer solutions.
For liquid water characteristic self-ionization . Its molecules mutually influence each other. The thermal movement of particles causes weakening and heterolytic rupture of O - H bonds in individual water molecules.
Ionic product of water– the product of concentrations [H + ] and – is a constant value at a constant temperature and equal to 10 -14 at 22°C.
The ionic product of water increases with increasing temperature.
pH value– negative logarithm of the concentration of hydrogen ions: pH = – log. Similarly: pOH = – log. Taking the logarithm of the ionic product of water gives: pH + pH = 14. The pH value characterizes the reaction of the medium. If pH = 7, then [H + ] = is a neutral medium.
If pH< 7, то [Н + ] >– acidic environment. If pH > 7, then [H + ]< – щелочная среда.
Buffer solutions– solutions having a certain concentration of hydrogen ions. The pH of these solutions does not change when diluted and changes little when small amounts of acids and alkalis are added.
The pH value of the solution is determined using a universal indicator.
Universal indicator is a mixture of several indicators that change color over a wide range of pH values.
37. Solubility of sparingly soluble solid electrolytes in water. Product of solubility (pr). Effect of common ions on solubility. Amphoteric hydroxides and oxides.
Solubility of a sparingly soluble substances can be expressed in moles per liter. Depending on the size s substances can be divided into poorly soluble – s< 10 -4 моль/л, среднерастворимые – 10 -4 моль/л ≤ s≤ 10 -2 mol/l and highly soluble s >10 -2 mol/l.
The solubility of compounds is related to their solubility product.
Solubility product (ETC, K sp) is the product of the concentration of ions of a slightly soluble electrolyte in its saturated solution at constant temperature and pressure. The solubility product is a constant value.
When introduced into sat. solution of a sparingly soluble electrolyte with a common ion, solubility decreases.
Amphoteric hydroxides- chemical substances that behave as bases in an acidic environment and as acids in an alkaline environment.
Amphoteric hydroxides practically insoluble in water, the most convenient way to obtain them is precipitation from an aqueous solution using a weak base - ammonia hydrate: Al(NO 3) 3 + 3(NH 3 H 2 O) = Al(OH) 3 ↓ + 3NH 4 NO 3 (20 °C) Al(NO 3) 3 + 3(NH 3 H 2 O) = AlO(OH)↓ + 3NH 4 NO 3 + H 2 O (80 °C)
Amphoteric oxides- salt-forming oxides that exhibit, depending on conditions, either basic or acidic properties (i.e.
exhibiting amphoteric properties). Formed by transition metals. Metals in amphoteric oxides usually exhibit valence II, III, IV.
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Electrolytic dissociation- this is the process of decomposition of a substance (which is an electrolyte) usually in water into ions that can move freely.
Acids in aqueous solutions are capable of dissociating into positively charged hydrogen ions (H+) and negatively charged acidic residues (for example, Cl -, SO 4 2-, NO 3 -). The former are called cations, the latter - anions. The sour taste of solutions of all acids is due to hydrogen ions.
Water molecules are polar. With their negatively charged poles they attract the hydrogen atoms of the acid, while other water molecules attract the acidic residues with their positively charged poles. If in an acid molecule the bond between hydrogen and the acidic residue is not strong enough, then it is broken, while the electron of the hydrogen atom remains with the acidic residue.
In solutions of strong acids, almost all molecules dissociate into ions. In weak acids, dissociation occurs more weakly, and along with it, the reverse process occurs - association - when the ions of the acidic residue and hydrogen form a bond, and again an electrically neutral acid molecule is obtained. Therefore, in dissociation equations, an equal sign or a unidirectional arrow is often used for strong acids, and multidirectional arrows are used for weak acids, thereby emphasizing that the process goes in both directions.
Strong electrolytes include hydrochloric (HCl), sulfuric (H 2 SO 4), nitric (HNO 3) acids, etc. Weak electrolytes include phosphoric (H 3 PO 4), nitrous (HNO 2), silicon (H 2 SiO 3) and etc.
A monobasic acid molecule (HCl, HNO 3, HNO 2, etc.) can dissociate into only one hydrogen ion and one acid residue ion. Thus, their dissociation always occurs in one step.
Molecules of polybasic acids (H 2 SO 4, H 3 PO 4, etc.) can dissociate in several stages. First, one hydrogen ion is split off from them, leaving a hydro-anion (for example, HSO 4 - hydro-sulfate ion). This is the first stage of dissociation. Next, the second hydrogen ion can be split off, leaving only an acidic residue (SO 4 2-). This is the second stage of dissociation.
Thus, the number of stages of electrolytic dissociation depends on the basicity of the acid (the number of hydrogen atoms in it).
The easiest way for dissociation to proceed is through the first step. With each subsequent step, dissociation decreases. The reason for this is that it is easier to remove a positively charged hydrogen ion from a neutral molecule than from a negatively charged one. After the first step, the remaining hydrogen ions are more strongly attracted to the acidic residue, since it has a greater negative charge.
By analogy with acids, bases also dissociate into ions. In this case, metal cations and hydroxide anions (OH -) are formed. Depending on the number of hydroxide groups in the base molecules, dissociation can also occur in several steps.
Polybasic acids and polyacid bases dissociate stepwise. Each dissociation step has its own dissociation constant. For example, for the dissociation of phosphoric acid:
The decrease in the constant from the first stage to the third is due to the fact that it becomes increasingly difficult to remove a proton as the negative charge of the resulting particle increases.
The total dissociation constant is equal to the product of the constants corresponding to the individual stages of dissociation. For example, in the case of phosphoric acid for the process:
To assess the degree of dissociation of weak electrolytes, it is sufficient to take into account only the first stage of dissociation – it, first of all, determines the concentration of ions in the solution.
Acidic and basic salts also dissociate in steps, for example:
It is easy to notice that the dissociation of a hydroanion or hydroxocation is identical to the second or third stage of dissociation of the corresponding acid or base and therefore obeys the same laws that have been formulated for the stepwise dissociation of acids and bases. In particular, if the basic salt corresponds to a weak base, and the acid salt – weak acid, then the dissociation of the hydroanion or hydroxocation (i.e., the second or third stage of salt dissociation) occurs to an insignificant extent.
Every oxygen-containing acid and every base (meaning acids and bases in the traditional sense) contain hydroxo groups. The difference between an acid and a base is that in the first case, dissociation occurs at the EO-H bond, and in the second – via E-ON connection.
Amphoteric hydroxides dissociate both as bases and as acids (both are very weak). Thus, the ionization of zinc hydroxide can be represented by the following scheme (without taking into account the hydration of the resulting ions):
The addition of acid shifts these equilibria to the left, and the addition of alkali – to the right. Therefore, in an acidic environment, dissociation according to the type of base predominates, and in an alkaline environment – by type of acid. In both cases, the binding of ions formed during the dissociation of a poorly soluble amphoteric electrolyte into water molecules causes the transition of new portions of such ions into the solution, their binding, the transition of new ions into the solution, etc. Consequently, the dissolution of such an electrolyte occurs both in an acid solution and and in an alkali solution.
(1887) to explain the properties of aqueous solutions of electrolytes. Subsequently, it was developed by many scientists on the basis of the doctrine of the structure of the atom and chemical bonds. The modern content of this theory can be reduced to the following three provisions:
Scheme for dissolving a crystal of table salt. Sodium and chlorine ions in solution.
1. Electrolytes, when dissolved in water, dissociate (break up) into ions - positively and negatively charged. (“Ion” is Greek for “wandering.” In a solution, ions move randomly in different directions.)
2. Under the influence of electric current, ions acquire directional movement: positively charged ones move towards the cathode, negatively charged ones move towards the anode. Therefore, the former are called cations, the latter - anions. The directional movement of ions occurs as a result of the attraction of their oppositely charged electrodes.
3. Dissociation is a reversible process. This means that a state of equilibrium occurs in which as many molecules break up into ions (dissociation), so many of them are formed again from ions (association). Therefore, in the equations of electrolytic dissociation, instead of the equal sign, the reversibility sign is used.
For example:
KA ↔ K + + A − ,
where KA is an electrolyte molecule, K + is a cation, A − is an anion.
The doctrine of chemical bonding helps answer the question of why electrolytes dissociate into ions. Substances with ionic bonds dissociate most easily, since they already consist of ions (see Chemical bonding). When they dissolve, the water dipoles are oriented around the positive and negative ions. Mutual attractive forces arise between the ions and dipoles of water. As a result, the bond between the ions weakens, and the ions move from the crystal to the solution. Electrolytes, whose molecules are formed according to the type of covalent polar bond, dissociate similarly. The dissociation of polar molecules can be complete or partial - it all depends on the degree of polarity of the bonds. In both cases (during the dissociation of compounds with ionic and polar bonds), hydrated ions are formed, that is, ions chemically bonded to water molecules.
The founder of this view of electrolytic dissociation was honorary academician I. A. Kablukov. In contrast to the Arrhenius theory, which did not take into account the interaction of the solute with the solvent, I. A. Kablukov applied the chemical theory of solutions of D. I. Mendeleev to explain electrolytic dissociation. He showed that during dissolution, a chemical interaction of the solute with water occurs, which leads to the formation of hydrates, and then they dissociate into ions. I. A. Kablukov believed that an aqueous solution contains only hydrated ions. Currently, this idea is generally accepted. So, ion hydration is the main cause of dissociation. In other, non-aqueous electrolyte solutions, the chemical bond between the particles (molecules, ions) of the solute and the solvent particles is called solvation.
Hydrated ions have both a constant and variable number of water molecules. A hydrate of constant composition forms hydrogen ions H + that hold one molecule of water - this is a hydrated proton H + (H 2 O). In the scientific literature, it is usually represented by the formula H 3 O + (or OH 3 +) and called the hydronium ion.
Since electrolytic dissociation is a reversible process, in solutions of electrolytes, along with their ions, there are also molecules. Therefore, electrolyte solutions are characterized by the degree of dissociation (denoted by the Greek letter a). The degree of dissociation is the ratio of the number of molecules disintegrated into ions, n, to the total number of dissolved molecules N:
The degree of electrolyte dissociation is determined experimentally and is expressed in fractions of a unit or as a percentage. If α = 0, then there is no dissociation, and if α = 1, or 100%, then the electrolyte completely disintegrates into ions. Different electrolytes have different degrees of dissociation. With dilution of the solution it increases, and with the addition of ions of the same name (the same as the electrolyte ions) it decreases.
However, to characterize the ability of an electrolyte to dissociate into ions, the degree of dissociation is not a very convenient value, since it... depends on the electrolyte concentration. A more general characteristic is the dissociation constant K. It can be easily derived by applying the law of mass action to the electrolyte dissociation equilibrium (1):
K = () / ,
where KA is the equilibrium concentration of the electrolyte, and are the equilibrium concentrations of its ions (see Chemical equilibrium). K does not depend on concentration. It depends on the nature of the electrolyte, solvent and temperature. For weak electrolytes, the higher K (dissociation constant), the stronger the electrolyte, the more ions in the solution.
Strong electrolytes do not have dissociation constants. Formally, they can be calculated, but they will not be constant as the concentration changes.
The theory of electrolytic dissociation proposed by the Swedish scientist S. Arrhenius in 1887.
Electrolytic dissociation- this is the breakdown of electrolyte molecules with the formation of positively charged (cations) and negatively charged (anions) ions in the solution.
For example, acetic acid dissociates like this in an aqueous solution:
CH 3 COOH⇄H + +CH 3 COO - .
Dissociation is a reversible process. But different electrolytes dissociate differently. The degree depends on the nature of the electrolyte, its concentration, the nature of the solvent, external conditions (temperature, pressure).
Dissociation degree α - the ratio of the number of molecules disintegrated into ions to the total number of molecules:
α=v´(x)/v(x).
The degree can vary from 0 to 1 (from no dissociation to its complete completion). Indicated as a percentage. Determined experimentally. When the electrolyte dissociates, the number of particles in the solution increases. The degree of dissociation indicates the strength of the electrolyte.
Distinguish strong And weak electrolytes.
Strong electrolytes- these are those electrolytes whose degree of dissociation exceeds 30%.
Medium strength electrolytes- these are those whose degree of dissociation ranges from 3% to 30%.
Weak electrolytes- the degree of dissociation in an aqueous 0.1 M solution is less than 3%.
Examples of weak and strong electrolytes.
Strong electrolytes in dilute solutions completely disintegrate into ions, i.e. α = 1. But experiments show that dissociation cannot be equal to 1, it has an approximate value, but is not equal to 1. This is not a true dissociation, but an apparent one.
For example, let some connection α = 0.7. Those. according to the Arrhenius theory, 30% of undissociated molecules “float” in the solution. And 70% formed free ions. And the electrostatic theory gives another definition to this concept: if α = 0.7, then all molecules are dissociated into ions, but the ions are only 70% free, and the remaining 30% are bound by electrostatic interactions.
Apparent degree of dissociation.
The degree of dissociation depends not only on the nature of the solvent and solute, but also on the concentration of the solution and temperature.
The dissociation equation can be represented as follows:
AK ⇄ A- + K + .
And the degree of dissociation can be expressed as follows:
As the solution concentration increases, the degree of electrolyte dissociation decreases. Those. the degree value for a particular electrolyte is not a constant value.
Since dissociation is a reversible process, the reaction rate equations can be written as follows:
If the dissociation is equilibrium, then the rates are equal and as a result we get equilibrium constant(dissociation constant):
K depends on the nature of the solvent and the temperature, but does not depend on the concentration of the solutions. It is clear from the equation that the more undissociated molecules there are, the lower the value of the electrolyte dissociation constant.
Polybasic acids dissociate stepwise, and each step has its own dissociation constant value.
If a polybasic acid dissociates, then the first proton is most easily removed, but as the charge of the anion increases, the attraction increases, and therefore the proton is much more difficult to remove. For example,
The dissociation constants of orthophosphoric acid at each step should vary greatly:
I - stage:
II - stage:
III - stage:
At the first stage, orthophosphoric acid is an acid of medium strength, and at the 2nd stage it is weak, at the 3rd stage it is very weak.
Examples of equilibrium constants for some electrolyte solutions.
Let's look at an example:
If metallic copper is added to a solution containing silver ions, then at the moment of equilibrium, the concentration of copper ions should be greater than the concentration of silver.
But the constant has a low value:
AgCl⇄Ag + +Cl - .
Which suggests that by the time equilibrium was reached, very little silver chloride had dissolved.
The metallic copper and silver concentrations are included in the equilibrium constant.
Ionic product of water.
The table below contains the following data:
This constant is called ionic product of water, which depends only on temperature. According to dissociation, there is one hydroxide ion per 1 H+ ion. In pure water the concentration of these ions is the same: [ H + ] = [OH - ].
From here, [ H + ] = [OH- ] = = 10-7 mol/l.
If you add a foreign substance, for example, hydrochloric acid, to water, the concentration of hydrogen ions will increase, but the ionic product of water does not depend on the concentration.
And if you add alkali, the concentration of ions will increase, and the amount of hydrogen will decrease.
Concentration and are interrelated: the greater one value, the less the other.
Acidity of the solution (pH).
The acidity of solutions is usually expressed by the concentration of ions H+. In acidic environments pH<10 -7 моль/л, в нейтральных - pH= 10 -7 mol/l, in alkaline - pH> 10 -7 mol/l.
The acidity of a solution is expressed through the negative logarithm of the concentration of hydrogen ions, calling it pH.
pH = -lg[ H + ].
The relationship between the constant and the degree of dissociation.
Consider an example of the dissociation of acetic acid:
Let's find the constant:
Molar concentration C=1/V, substitute it into the equation and get:
These equations are W. Ostwald's breeding law, according to which the dissociation constant of the electrolyte does not depend on the dilution of the solution.
