Which are in dynamic equilibrium with undissociated molecules. Weak electrolytes include most organic acids and many organic bases in aqueous and non-aqueous solutions.
Weak electrolytes are:
- almost all organic acids and water;
- some inorganic acids: HF, HClO, HClO 2, HNO 2, HCN, H 2 S, HBrO, H 3 PO 4, H 2 CO 3, H 2 SiO 3, H 2 SO 3, etc.;
- some poorly soluble metal hydroxides: Fe(OH) 3, Zn(OH) 2, etc.; as well as ammonium hydroxide NH 4 OH.
Literature
- M. I. Ravich-Sherbo. V.V. Novikov “Physical and colloidal Chemistry” M: graduate School, 1975
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See what “Weak electrolytes” are in other dictionaries:
weak electrolytes- – electrolytes that slightly dissociate into ions in aqueous solutions. Process of dissociation weak electrolytes reversible and obeys the law of mass action. general chemistry: textbook / A. V. Zholnin ... Chemical terms
Substances with ionic conductivity; They are called conductors of the second kind; the passage of current through them is accompanied by the transfer of matter. Electrolytes include molten salts, oxides or hydroxides, as well as (which occurs significantly... ... Collier's Encyclopedia
In a broad sense, liquid or solid systems in which ions are present in a noticeable concentration, causing the passage of electricity through them. current (ionic conductivity); in the narrow sense, in va, which disintegrate in p re into ions. When dissolving E.... ... Physical encyclopedia
Electrolytes- liquid or solids, in which, as a result of electrolytic dissociation, ions are formed in any noticeable concentration, causing the passage of a constant electric current. Electrolytes in solutions... ... Encyclopedic Dictionary of Metallurgy
In va, in which ions are present in noticeable concentrations, causing the passage of electricity. current (ionic conductivity). E. also called. conductors of the second kind. In the narrow sense of the word, E. in va, molecules that are in p re due to electrolytic ... ... Chemical encyclopedia
- (from Electro... and Greek lytos decomposed, soluble) liquid or solid substances and systems in which ions are present in any noticeable concentration, causing the passage of electric current. In the narrow sense, E.... ... Great Soviet Encyclopedia
This term has other meanings, see Dissociation. Electrolytic dissociation is the process of the breakdown of an electrolyte into ions when it dissolves or melts. Contents 1 Dissociation in solutions 2 ... Wikipedia
An electrolyte is a substance whose melt or solution conducts electric current due to dissociation into ions, but the substance itself does not conduct electric current. Examples of electrolytes are solutions of acids, salts and bases.... ... Wikipedia
Electrolyte is a chemical term denoting a substance whose melt or solution conducts electric current due to dissociation into ions. Examples of electrolytes include acids, salts and bases. Electrolytes are conductors of the second kind, ... ... Wikipedia
The value of a is expressed in fractions of a unit or in % and depends on the nature of the electrolyte, solvent, temperature, concentration and composition of the solution.
The solvent plays a special role: in some cases, when moving from aqueous solutions to organic solvents the degree of dissociation of electrolytes may sharply increase or decrease. In the following, in the absence of special instructions, we will assume that the solvent is water.
According to the degree of dissociation, electrolytes are conventionally divided into strong(a > 30%), average (3% < a < 30%) и weak(a< 3%).
Strong electrolytes include:
1) some inorganic acids (HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 4 and a number of others);
2) hydroxides of alkali (Li, Na, K, Rb, Cs) and alkaline earth (Ca, Sr, Ba) metals;
3) almost all soluble salts.
Electrolytes of medium strength include Mg(OH) 2, H 3 PO 4, HCOOH, H 2 SO 3, HF and some others.
All carboxylic acids (except HCOOH) and hydrated forms of aliphatic and aromatic amines are considered weak electrolytes. Many inorganic acids (HCN, H 2 S, H 2 CO 3, etc.) and bases (NH 3 ∙H 2 O) are also weak electrolytes.
Despite some similarities, in general one should not equate the solubility of a substance with its degree of dissociation. So, acetic acid and ethanol are unlimitedly soluble in water, but at the same time the first substance is a weak electrolyte, and the second is a non-electrolyte.
Acids and bases
Despite the fact that the concepts “acid” and “base” are widely used to describe chemical processes, there is no single approach to the classification of substances in terms of classifying them as acids or bases. Currently existing theories ( ionic theory S. Arrhenius, protolytic theory I. Brønsted and T. Lowry And electronic theory G. Lewis) have certain limitations and are therefore only applicable in special cases. Let's take a closer look at each of these theories.
Arrhenius theory.
In Arrhenius's ionic theory, the concepts of "acid" and "base" are closely related to the process of electrolytic dissociation:
An acid is an electrolyte that dissociates in solutions to form H + ions;
The base is an electrolyte that dissociates in solutions to form OH - ions;
An ampholyte (amphoteric electrolyte) is an electrolyte that dissociates in solutions to form both H + ions and OH - ions.
For example:
HA ⇄ H + + A - nH + + MeO n n - ⇄ Me(OH) n ⇄ Me n + + nOH -According to the ionic theory, acids can be either neutral molecules or ions, for example:
HF ⇄ H + + F -
H 2 PO 4 - ⇄ H + + HPO 4 2 -
NH 4 + ⇄ H + + NH 3
Similar examples can be given for grounds:
KOH K + + OH -
- ⇄ Al(OH) 3 + OH -
+ ⇄ Fe 2+ + OH -
Ampholytes include hydroxides of zinc, aluminum, chromium and some others, as well as amino acids, proteins, and nucleic acids.
In general, acid-base interaction in a solution comes down to a neutralization reaction:
H + + OH - H 2 O
However, a number of experimental data show the limitations of the ionic theory. So, ammonia, organic amines, metal oxides such as Na 2 O, CaO, anions of weak acids, etc. in the absence of water they exhibit the properties of typical bases, although they do not contain hydroxide ions.
On the other hand, many oxides (SO 2 , SO 3 , P 2 O 5 , etc.), halides, acid halides, without containing hydrogen ions, exhibit acidic properties even in the absence of water, i.e. neutralize bases.
In addition, the behavior of an electrolyte in an aqueous solution and in a non-aqueous medium may be opposite.
So, CH 3 COOH in water is a weak acid:
CH 3 COOH ⇄ CH 3 COO - + H + ,
and in liquid hydrogen fluoride it exhibits the properties of a base:
HF + CH 3 COOH ⇄ CH 3 COOH 2 + + F -
Studies of these types of reactions, and especially reactions occurring in non-aqueous solvents, have led to the development of more general theories acids and bases.
The theory of Bronsted and Lowry.
A further development of the theory of acids and bases was the protolytic (proton) theory proposed by I. Brønsted and T. Lowry. According to this theory:
An acid is any substance whose molecules (or ions) are capable of donating a proton, i.e. be a proton donor;
A base is any substance whose molecules (or ions) are capable of attaching a proton, i.e. be a proton acceptor;
Thus, the concept of foundation is significantly expanded, which is confirmed by the following reactions:
OH - + H + H 2 O
NH 3 + H + NH 4 +
H 2 N-NH 3 + + H + H 3 N + -NH 3 +
According to the theory of I. Brønsted and T. Lowry, an acid and a base form a conjugate pair and are related by equilibrium:
ACID ⇄ PROTON + BASE
Since the proton transfer reaction (protolytic reaction) is reversible, and a proton is also transferred in the reverse process, the reaction products are acids and bases in relation to each other. This can be written as an equilibrium process:
NA + B ⇄ VN + + A - ,
where HA is an acid, B is a base, BH + is an acid conjugate to base B, A - is a base conjugate to acid HA.
Examples.
1) in the reaction:
HCl + OH - ⇄ Cl - + H 2 O,
HCl and H 2 O are acids, Cl - and OH - are the corresponding bases conjugate to them;
2) in the reaction:
HSO 4 - + H 2 O ⇄ SO 4 2 - + H 3 O +,
HSO 4 - and H 3 O + are acids, SO 4 2 - and H 2 O are bases;
3) in the reaction:
NH 4 + + NH 2 - ⇄ 2NH 3,
NH 4 + is an acid, NH 2 - is a base, and NH 3 acts as both an acid (one molecule) and a base (another molecule), i.e. demonstrates signs of amphotericity - the ability to exhibit the properties of an acid and a base.
Water also has this ability:
2H 2 O ⇄ H 3 O + + OH -
Here, one molecule H 2 O attaches a proton (base), forming a conjugate acid - hydronium ion H 3 O +, the other gives up a proton (acid), forming a conjugate base OH -. This process is called autoprotolysis.
From the above examples it is clear that, in contrast to the ideas of Arrhenius, in the theory of Brønsted and Lowry, reactions of acids with bases do not lead to mutual neutralization, but are accompanied by the formation of new acids and bases.
It should also be noted that the protolytic theory considers the concepts of “acid” and “base” not as a property, but as a function that the compound in question performs in a protolytic reaction. The same compound can react as an acid under some conditions and as a base under others. Thus, in an aqueous solution, CH 3 COOH exhibits the properties of an acid, and in 100% H 2 SO 4 it exhibits the properties of a base.
However, despite its advantages, the protolytic theory, like the Arrhenius theory, is not applicable to substances that do not contain hydrogen atoms, but, at the same time, exhibit the function of an acid: boron, aluminum, silicon, tin halides.
Lewis' theory.
Another approach to the classification of substances from the point of view of classifying them as acids and bases was the Lewis electron theory. Within the framework of electronic theory:
an acid is a particle (molecule or ion) capable of attaching an electron pair (electron acceptor);
A base is a particle (molecule or ion) capable of donating an electron pair (electron donor).
According to Lewis's ideas, an acid and a base interact with each other to form a donor-acceptor bond. As a result of the addition of a pair of electrons to an atom with an electron deficiency, a complete electronic configuration- octet of electrons. For example:
The reaction between neutral molecules can be imagined in a similar way:
The neutralization reaction in terms of the Lewis theory is considered as the addition of an electron pair of a hydroxide ion to a hydrogen ion, which provides a free orbital to accommodate this pair:
Thus, the proton itself, which easily attaches an electron pair, from the point of view of the Lewis theory, performs the function of an acid. In this regard, Bronsted acids can be considered as reaction products between Lewis acids and bases. Thus, HCl is the product of neutralization of the acid H + with the base Cl -, and the H 3 O + ion is formed as a result of the neutralization of the acid H + with the base H 2 O.
Reactions between Lewis acids and bases are also illustrated by the following examples:
Lewis bases also include halide ions, ammonia, aliphatic and aromatic amines, oxygen-containing organic compounds such as R 2 CO (where R is an organic radical).
Lewis acids include halides of boron, aluminum, silicon, tin and other elements.
It is obvious that in Lewis's theory the concept of "acid" includes a wider range chemical compounds. This is explained by the fact that, according to Lewis, the classification of a substance as an acid is determined solely by the structure of its molecule, which determines the electron-acceptor properties, and is not necessarily related to the presence of hydrogen atoms. Lewis acids that do not contain hydrogen atoms are called aprotic.
Problem solving standards
1. Write the equation for the electrolytic dissociation of Al 2 (SO 4) 3 in water.
Aluminum sulfate is a strong electrolyte and in aqueous solution undergoes complete decomposition into ions. Dissociation equation:
Al 2 (SO 4) 3 + (2x + 3y)H 2 O 2 3+ + 3 2 - ,
or (without taking into account the process of ion hydration):
Al 2 (SO 4) 3 2Al 3+ + 3SO 4 2 - .
2. What is the HCO 3 ion from the perspective of the Brønsted-Lowry theory?
Depending on the conditions, the HCO 3 ion can donate protons:
HCO 3 - + OH - CO 3 2 - + H 2 O (1),
add protons like this:
HCO 3 - + H 3 O + H 2 CO 3 + H 2 O (2).
Thus, in the first case, the HCO 3 - ion is an acid, in the second, it is a base, i.e., it is an ampholyte.
3. Determine what the Ag + ion is in the reaction from the standpoint of Lewis theory:
Ag + + 2NH 3 +
In the process of formation of chemical bonds, which proceeds according to the donor-acceptor mechanism, the Ag + ion, having a free orbital, is an acceptor of electron pairs, and thus exhibits the properties of a Lewis acid.
4. Determine the ionic strength of a solution containing 0.1 mol KCl and 0.1 mol Na 2 SO 4 in one liter.
The dissociation of the presented electrolytes proceeds in accordance with the equations:
Na 2 SO 4 2Na + + SO 4 2 -
Hence: C(K +) = C(Cl -) = C(KCl) = 0.1 mol/l;
C(Na +) = 2×C(Na 2 SO 4) = 0.2 mol/l;
C(SO 4 2 -) = C(Na 2 SO 4) = 0.1 mol/l.
The ionic strength of the solution is calculated using the formula:
5. Determine the concentration of CuSO 4 in a solution of this electrolyte with I= 0.6 mol/l.
The dissociation of CuSO 4 proceeds according to the equation:
CuSO 4 Cu 2+ + SO 4 2 -
Let's take C(CuSO 4) as x mol/l, then, in accordance with the reaction equation, C(Cu 2+) = C(SO 4 2 -) = x mol/l. IN in this case the expression for calculating ionic strength will be:
6. Determine the activity coefficient of the K + ion in an aqueous solution of KCl with C(KCl) = 0.001 mol/l.
which in this case will take the form:
.
We find the ionic strength of the solution using the formula:
7. Determine the activity coefficient of the Fe 2+ ion in an aqueous solution whose ionic strength is 1.
According to the Debye-Hückel law:
hence:
8. Determine the dissociation constant of acid HA if in a solution of this acid with a concentration of 0.1 mol/l a = 24%.
Based on the degree of dissociation, it can be determined that this acid is an electrolyte of medium strength. Therefore, to calculate the acid dissociation constant, we use the Ostwald dilution law in its full form:
9. Determine the electrolyte concentration if a = 10%, K d = 10 - 4.
From Ostwald's law of dilution:
10. The degree of dissociation of monobasic acid HA does not exceed 1%. (HA) = 6.4×10 - 7. Determine the degree of dissociation of HA in its solution with a concentration of 0.01 mol/L.
Based on the degree of dissociation, it can be determined that this acid is a weak electrolyte. This allows us to use the approximate formula of Ostwald's dilution law:
11. The degree of dissociation of the electrolyte in its solution with a concentration of 0.001 mol/l is 0.009. Determine the dissociation constant of this electrolyte.
From the conditions of the problem it is clear that this electrolyte is weak (a = 0.9%). That's why:
12. (HNO 2) = 3.35. Compare the strength of HNO 2 with the strength of monobasic acid HA, the degree of dissociation of which in a solution with C(HA) = 0.15 mol/l is 15%.
Let's calculate (HA) using full form Ostwald equations:
Since (HA)< (HNO 2), то кислота HA является более сильной кислотой по сравнению с HNO 2 .
13. There are two solutions of KCl, which also contain other ions. It is known that the ionic strength of the first solution ( I 1) is equal to 1, and the second ( I 2) is 10 - 2 . Compare activity rates f(K +) in these solutions and conclude how the properties of these solutions differ from the properties of infinitely dilute KCl solutions.
We calculate the activity coefficients of K + ions using the Debye-Hückel law:
Activity factor f is a measure of the deviation in the behavior of an electrolyte solution of a given concentration from its behavior when the solution is infinitely diluted.
Because f 1 = 0.316 deviates more from 1 than f 2 = 0.891, then in a solution with higher ionic strength there is a greater deviation in the behavior of the KCl solution from its behavior at infinite dilution.
Questions for self-control
1. What is electrolytic dissociation?
2. What substances are called electrolytes and non-electrolytes? Give examples.
3. What is the degree of dissociation?
4. On what factors does the degree of dissociation depend?
5. Which electrolytes are considered strong? Which are medium strength? Which ones are weak? Give examples.
6. What is the dissociation constant? What does the dissociation constant depend on and what does it not depend on?
7. How are the constant and the degree of dissociation related to each other in binary solutions of medium and weak electrolytes?
8. Why do solutions of strong electrolytes show deviations from ideality in their behavior?
9. What is the meaning of the term “apparent degree of dissociation”?
10. What is the activity of an ion? What is the activity coefficient?
11. How does the activity coefficient change with dilution (concentration) of a strong electrolyte solution? What is the limiting value of the activity coefficient for an infinite solution dilution?
12. What is the ionic strength of a solution?
13. How is the activity coefficient calculated? Formulate the Debye-Hückel law.
14. What is the essence of the ionic theory of acids and bases (Arrhenius theory)?
15. What is it? fundamental difference protolytic theory of acids and bases (theory of Brønsted and Lowry) from the theory of Arrhenius?
16. How does electronic theory (Lewis theory) interpret the concepts of “acid” and “base”? Give examples.
Task options for independent decision
Option #1
1. Write the equation for the electrolytic dissociation of Fe 2 (SO 4) 3.
HA + H 2 O ⇄ H 3 O + + A - .
Option No. 2
1. Write the equation for the electrolytic dissociation of CuCl 2.
2. Determine what the S 2 - ion is in the reaction from the standpoint of Lewis theory:
2Ag + + S 2 - ⇄ Ag 2 S.
3. Calculate the molar concentration of the electrolyte in the solution if a = 0.75%, a = 10 - 5.
Option #3
1. Write the equation for the electrolytic dissociation of Na 2 SO 4.
2. Determine what the CN - ion is in the reaction from the standpoint of Lewis theory:
Fe 3 + + 6CN - ⇄ 3 - .
3. The ionic strength of the CaCl 2 solution is 0.3 mol/l. Calculate C(CaCl2).
Option No. 4
1. Write the equation for the electrolytic dissociation of Ca(OH) 2.
2. Determine what the H 2 O molecule is in the reaction from the standpoint of the Brønsted theory:
H 3 O + ⇄ H + + H 2 O.
3. The ionic strength of the K 2 SO 4 solution is 1.2 mol/L. Calculate C(K 2 SO 4).
Option #5
1. Write the equation for the electrolytic dissociation of K 2 SO 3.
NH 4 + + H 2 O ⇄ NH 3 + H 3 O + .
3. (CH 3 COOH) = 4.74. Compare the strength of CH 3 COOH with the strength of monobasic acid HA, the degree of dissociation of which in solution with C(HA) = 3.6 × 10 - 5 mol/l is 10%.
Option #6
1. Write the equation for electrolytic dissociation of K 2 S.
2. Determine what the AlBr 3 molecule is in the reaction from the standpoint of Lewis theory:
Br - + AlBr 3 ⇄ - .
Option No. 7
1. Write the equation for the electrolytic dissociation of Fe(NO 3) 2.
2. Determine what the Cl - ion is in the reaction from the standpoint of Lewis theory:
Cl - + AlCl 3 ⇄ - .
Option No. 8
1. Write the equation for the electrolytic dissociation of K 2 MnO 4 .
2. Determine what the HSO 3 - ion is in the reaction from the standpoint of the Brønsted theory:
HSO 3 - + OH – ⇄ SO 3 2 - + H 2 O.
Option No. 9
1. Write the equation for the electrolytic dissociation of Al 2 (SO 4) 3.
2. Determine what the Co 3+ ion is in the reaction from the standpoint of Lewis theory:
Co 3+ + 6NO 2 - ⇄ 3 - .
3. 1 liter of solution contains 0.348 g of K2SO4 and 0.17 g of NaNO3. Determine the ionic strength of this solution.
Option No. 10
1. Write the equation for the electrolytic dissociation of Ca(NO 3) 2.
2. Determine what the H 2 O molecule is in the reaction from the standpoint of the Brønsted theory:
B + H 2 O ⇄ OH - + BH + .
3. Calculate the electrolyte concentration in the solution if a = 5%, a = 10 - 5.
Option No. 11
1. Write the equation for the electrolytic dissociation of KMnO 4.
2. Determine what the Cu 2+ ion is in the reaction from the perspective of Lewis theory:
Cu 2+ + 4NH 3 ⇄ 2 + .
3. Calculate the activity coefficient of the Cu 2+ ion in a solution of CuSO 4 with C(CuSO 4) = 0.016 mol/l.
Option No. 12
1. Write the equation for the electrolytic dissociation of Na 2 CO 3.
2. Determine what the H 2 O molecule is in the reaction from the standpoint of the Brønsted theory:
K + + xH 2 O ⇄ + .
3. There are two NaCl solutions containing other electrolytes. The ionic strengths of these solutions are respectively equal: I 1 = 0.1 mol/l, I 2 = 0.01 mol/l. Compare activity rates f(Na +) in these solutions.
Option No. 13
1. Write the equation for the electrolytic dissociation of Al(NO 3) 3.
2. Determine what the RNH 2 molecule is in the reaction from the standpoint of Lewis theory:
RNH 2 + H 3 O + ⇄ RNH 3 + + H 2 O.
3. Compare the activity coefficients of cations in a solution containing FeSO 4 and KNO 3, provided that the electrolyte concentrations are 0.3 and 0.1 mol/l, respectively.
Option No. 14
1. Write the equation for the electrolytic dissociation of K 3 PO 4.
2. Determine what the H 3 O + ion is in the reaction from the standpoint of the Brønsted theory:
HSO 3 - + H 3 O + ⇄ H 2 SO 3 + H 2 O.
Option No. 15
1. Write the equation for the electrolytic dissociation of K 2 SO 4.
2. Determine what Pb(OH) 2 is in the reaction from the standpoint of Lewis theory:
Pb(OH) 2 + 2OH - ⇄ 2 - .
Option No. 16
1. Write the equation for the electrolytic dissociation of Ni(NO 3) 2.
2. Determine what the hydronium ion (H 3 O +) is in the reaction from the standpoint of the Brønsted theory:
2H 3 O + + S 2 - ⇄ H 2 S + 2H 2 O.
3. The ionic strength of a solution containing only Na 3 PO 4 is 1.2 mol/l. Determine the concentration of Na 3 PO 4.
Option No. 17
1. Write the equation for the electrolytic dissociation of (NH 4) 2 SO 4.
2. Determine what the NH 4 + ion is in the reaction from the standpoint of the Brønsted theory:
NH 4 + + OH - ⇄ NH 3 + H 2 O.
3. The ionic strength of a solution containing both KI and Na 2 SO 4 is 0.4 mol/l. C(KI) = 0.1 mol/l. Determine the concentration of Na 2 SO 4.
Option No. 18
1. Write the equation for the electrolytic dissociation of Cr 2 (SO 4) 3.
2. Determine what the protein molecule in the reaction is from the perspective of the Brønsted theory:
INFORMATION BLOCK
pH scale
Table 3. Relationship between the concentrations of H + and OH - ions.
Problem solving standards
1. The concentration of hydrogen ions in the solution is 10 - 3 mol/l. Calculate the pH, pOH and [OH - ] values in this solution. Determine the solution medium.
Note. The following ratios are used for calculations: lg10 a = a; 10 lg a = A.
A solution environment with pH = 3 is acidic, since the pH< 7.
2. Calculate the pH of the solution of hydrochloric acid with a molar concentration of 0.002 mol/l.
Since in a dilute solution HC1 » 1, and in a solution of a monobasic acid C(s) = C(s), we can write:
3. 90 ml of water was added to 10 ml of acetic acid solution with C(CH 3 COOH) = 0.01 mol/l. Find the difference in pH values of the solution before and after dilution, if (CH 3 COOH) = 1.85 × 10 - 5.
1) In the initial solution of a weak monobasic acid CH 3 COOH:
Hence:
2) Adding 90 ml of water to 10 ml of acid solution corresponds to a 10-fold dilution of the solution. That's why.
Electrolytes are substances, alloys of substances or solutions that have the ability to electrolytically conduct galvanic current. You can determine which electrolytes a substance belongs to using the theory of electrolytic dissociation.
Instructions
- The essence of this theory is that when melted (dissolved in water), almost all electrolytes are decomposed into ions, which are both positively and negatively charged (which is called electrolytic dissociation). Under the influence of electric current, negative ones (anions, “-”) move towards the anode (+), and positively charged ones (cations, “+”) move towards the cathode (-). Electrolytic dissociation is a reversible process ( reverse process is called "molarization").
- The degree of (a) electrolytic dissociation depends on the nature of the electrolyte itself, the solvent, and their concentration. This is the ratio of the number of molecules (n) that have broken up into ions to total number molecules (N) introduced into the solution. You get: a = n / N
- Thus, strong electrolytes are substances that completely disintegrate into ions when dissolved in water. Strong electrolytes, as a rule, include substances with highly polar or ionic bonds: these are salts that are highly soluble, strong acids (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 (anions and cations); 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. Weak electrolytes include:
- organic acids (almost all) (C2H5COOH, CH3COOH, etc.);
- some of the inorganic 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.
Salts, their properties, hydrolysis
8th grade student B of school No. 182
Petrova Polina
Chemistry teacher:
Kharina Ekaterina Alekseevna
MOSCOW 2009
In everyday life, we are accustomed to dealing with only one salt - table salt, i.e. sodium chloride NaCl. However, in chemistry, a whole class of compounds is called salts. Salts can be considered as products of the replacement of hydrogen in an acid with a metal. Table salt, for example, can be obtained from hydrochloric acid by a substitution reaction:
2Na + 2HCl = 2NaCl + H2.
acid salt
If you take aluminum instead of sodium, another salt is formed - aluminum chloride:
2Al + 6HCl = 2AlCl3 + 3H2
Salts- These are complex substances consisting of metal atoms and acidic residues. They are products of complete or partial replacement of hydrogen in an acid with a metal or hydroxyl group in the base. acid residue. For example, if in sulfuric acid H 2 SO 4 we replace one hydrogen atom with potassium, we get the salt KHSO 4, and if two - K 2 SO 4.
There are several types of salts.
| Types of salts | Definition | Examples of salts |
| Average | The product of complete replacement of acid hydrogen with metal. They contain neither H atoms nor OH groups. | Na 2 SO 4 sodium sulfate CuCl 2 copper (II) chloride Ca 3 (PO 4) 2 calcium phosphate Na 2 CO 3 sodium carbonate (soda ash) |
| Sour | A product of incomplete replacement of acid hydrogen by metal. Contain hydrogen atoms. (They are formed only by polybasic acids) | CaHPO 4 calcium hydrogen phosphate Ca(H 2 PO 4) 2 calcium dihydrogen phosphate NaHCO 3 sodium bicarbonate (baking soda) |
| Basic | The product of incomplete replacement of the hydroxyl groups of a base with an acidic residue. Includes OH groups. (Formed only by polyacid bases) | Cu(OH)Cl copper (II) hydroxychloride Ca 5 (PO 4) 3 (OH) calcium hydroxyphosphate (CuOH) 2 CO 3 copper (II) hydroxycarbonate (malachite) |
| Mixed | Salts of two acids | Ca(OCl)Cl – bleach |
| Double | Salts of two metals | K 2 NaPO 4 – dipotassium sodium orthophosphate |
| Crystalline hydrates | Contains water of crystallization. When heated, they dehydrate - they lose water, turning into anhydrous salt. | CuSO4. 5H 2 O – copper(II) sulfate pentahydrate ( copper sulfate) Na 2 CO 3 . 10H 2 O – sodium carbonate decahydrate (soda) |
Methods for obtaining salts.
1. Salts can be obtained by acting with acids on metals, basic oxides and bases:
Zn + 2HCl ZnCl 2 + H 2
zinc chloride
3H 2 SO 4 + Fe 2 O 3 Fe 2 (SO 4) 3 + 3H 2 O
iron(III) sulfate
3HNO 3 + Cr(OH) 3 Cr(NO 3) 3 + 3H 2 O
chromium(III) nitrate
2. Salts are formed by the reaction of acidic oxides with alkalis, as well as acidic oxides with basic oxides:
N 2 O 5 + Ca(OH) 2 Ca(NO 3) 2 + H 2 O
calcium nitrate
SiO 2 + CaO CaSiO 3
calcium silicate
3. Salts can be obtained by reacting salts with acids, alkalis, metals, non-volatile acid oxides and other salts. Such reactions occur under the conditions of evolution of gas, precipitation of a precipitate, evolution of an oxide of a weaker acid, or evolution of a volatile oxide.
Ca 3 (PO4) 2 + 3H 2 SO 4 3CaSO 4 + 2H 3 PO 4
calcium orthophosphate calcium sulfate
Fe 2 (SO 4) 3 + 6NaOH 2Fe(OH) 3 + 3Na 2 SO 4
iron (III) sulfate sodium sulfate
CuSO 4 + Fe FeSO 4 + Cu
copper (II) sulfate iron (II) sulfate
CaCO 3 + SiO 2 CaSiO 3 + CO 2
calcium carbonate calcium silicate
Al 2 (SO 4) 3 + 3BaCl 2 3BaSO 4 + 2AlCl 3
sulfate chloride sulfate chloride
aluminum barium barium aluminum
4. Salts of oxygen-free acids are formed by the interaction of metals with non-metals:
2Fe + 3Cl 2 2FeCl 3
iron(III) chloride
Physical properties.
Salts are solids different colors. Their solubility in water varies. All salts of nitric and acetic acids, as well as sodium and potassium salts, are soluble. The solubility of other salts in water can be found in the solubility table.
Chemical properties.
1) Salts react with metals.
Since these reactions occur in aqueous solutions, Li, Na, K, Ca, Ba and other active metals that react with water under normal conditions cannot be used for experiments, or reactions cannot be carried out in a melt.
CuSO 4 + Zn ZnSO 4 + Cu
Pb(NO 3) 2 + Zn Zn(NO 3) 2 + Pb
2) Salts react with acids. These reactions occur when a stronger acid displaces a weaker one, releasing gas or precipitating.
When carrying out these reactions, they usually take dry salt and act with concentrated acid.
BaCl 2 + H 2 SO 4 BaSO 4 + 2HCl
Na 2 SiO 3 + 2HCl 2NaCl + H 2 SiO 3
3) Salts react with alkalis in aqueous solutions.
This is a method of obtaining insoluble bases and alkalis.
FeCl 3 (p-p) + 3NaOH(p-p) Fe(OH) 3 + 3NaCl
CuSO 4 (p-p) + 2NaOH (p-p) Na 2 SO 4 + Cu(OH) 2
Na 2 SO 4 + Ba(OH) 2 BaSO 4 + 2NaOH
4) Salts react with salts.
The reactions take place in solutions and are used to obtain practically insoluble salts.
AgNO 3 + KBr AgBr + KNO 3
CaCl 2 + Na 2 CO 3 CaCO 3 + 2NaCl
5) Some salts decompose when heated.
A typical example of such a reaction is the firing of limestone, the main integral part which is calcium carbonate:
CaCO 3 CaO + CO2 calcium carbonate
1. Some salts are capable of crystallizing to form crystalline hydrates.
Copper (II) sulfate CuSO 4 – crystalline substance white. When it dissolves in water, it heats up and forms a solution blue color. Warmth and color changes are signs chemical reaction. When the solution is evaporated, crystalline hydrate CuSO 4 is released. 5H 2 O (copper sulfate). The formation of this substance indicates that copper (II) sulfate reacts with water:
CuSO 4 + 5H 2 O CuSO 4 . 5H 2 O + Q
white blue-blue color
Use of salts.
Most salts are widely used in industry and in everyday life. For example, sodium chloride NaCl, or table salt, is indispensable in cooking. In industry, sodium chloride is used to produce sodium hydroxide, soda NaHCO 3, chlorine, sodium. Salts of nitric and orthophosphoric acids are mainly mineral fertilizers. For example, potassium nitrate KNO 3 is potassium nitrate. It is also part of gunpowder and other pyrotechnic mixtures. Salts are used to obtain metals, acids, and in glass production. Many plant protection products from diseases, pests, and some medicinal substances also belong to the class of salts. Potassium permanganate KMnO 4 is often called potassium permanganate. As building material limestone and gypsum are used - CaSO 4. 2H 2 O, which is also used in medicine.
Solutions and solubility.
As stated earlier, solubility is important property salts Solubility is the ability of a substance to form a homogeneous substance with another substance, sustainable system variable composition, consisting of two or more components.
Solutions- This homogeneous systems, consisting of solvent molecules and solute particles.
So, for example, a solution table salt consists of a solvent - water, a dissolved substance - Na +, Cl - ions.
Ions(from Greek ión - going), electrically charged particles formed by the loss or gain of electrons (or other charged particles) by atoms or groups of atoms. The concept and term “ion” was introduced in 1834 by M. Faraday, who, while studying the effect of electric current on aqueous solutions of acids, alkalis and salts, suggested that the electrical conductivity of such solutions is due to the movement of ions. Faraday called positively charged ions moving in solution towards the negative pole (cathode) cations, and negatively charged ions moving towards the positive pole (anode) - anions.
Based on the degree of solubility in water, substances are divided into three groups:
1) Highly soluble;
2) Slightly soluble;
3) Practically insoluble.
Many salts are highly soluble in water. When deciding the solubility of other salts in water, you will have to use the solubility table.
It is well known that some substances, when dissolved or molten, conduct electric current, while others do not conduct current under the same conditions.
Substances that disintegrate into ions in solutions or melts and therefore conduct electric current are called electrolytes.
Substances that, under the same conditions, do not disintegrate into ions and do not conduct electric current are called non-electrolytes.
Electrolytes include acids, bases and almost all salts. Electrolytes themselves do not conduct electricity. In solutions and melts, they break up into ions, which is why current flows.
The breakdown of electrolytes into ions when dissolved in water is called electrolytic dissociation. Its content boils down to the following three provisions:
1) Electrolytes, when dissolved in water, break up (dissociate) into ions - positive and negative.
2) Under the influence of an electric current, ions acquire directional movement: positively charged ions move towards the cathode and are called cations, and negatively charged ions move towards the anode and are called anions.
3) Dissociation is a reversible process: in parallel with the disintegration of molecules into ions (dissociation), the process of combining ions (association) occurs.
reversibility
Strong and weak electrolytes.
To quantitatively characterize the ability of an electrolyte to disintegrate into ions, the concept of the degree of dissociation (α), t . E. The ratio of the number of molecules disintegrated into ions to the total number of molecules. For example, α = 1 indicates that the electrolyte has completely disintegrated into ions, and α = 0.2 means that only every fifth of its molecules has dissociated. When a concentrated solution is diluted, as well as when heated, its electrical conductivity increases, as the degree of dissociation increases.
Depending on the value of α, electrolytes are conventionally divided into strong (dissociate almost completely, (α 0.95)) medium strength (0.95
Strong electrolytes are many mineral acids (HCl, HBr, HI, H 2 SO 4, HNO 3, etc.), alkalis (NaOH, KOH, Ca(OH) 2, etc.), and almost all salts. Weak ones include solutions of some mineral acids (H 2 S, H 2 SO 3, H 2 CO 3, HCN, HClO), many organic acids (for example, acetic acid CH 3 COOH), an aqueous solution of ammonia (NH 3. 2 O), water, some mercury salts (HgCl 2). Electrolytes of medium strength often include hydrofluoric HF, orthophosphoric H 3 PO 4 and nitrous HNO 2 acids.
Hydrolysis of salts.
The term "hydrolysis" comes from Greek words hidor (water) and lysis (decomposition). Hydrolysis is usually understood as an exchange reaction between a substance and water. Hydrolytic processes are extremely common in the nature around us (both living and nonliving), and are also widely used by humans in modern production and household technologies.
Salt hydrolysis is the reaction of interaction between the ions that make up the salt and water, which leads to the formation of a weak electrolyte and is accompanied by a change in the solution environment.
Three types of salts undergo hydrolysis:
a) salts formed by a weak base and a strong acid (CuCl 2, NH 4 Cl, Fe 2 (SO 4) 3 - hydrolysis of the cation occurs)
NH 4 + + H 2 O NH 3 + H 3 O +
NH 4 Cl + H 2 O NH 3 . H2O + HCl
The reaction of the medium is acidic.
b) salts formed by a strong base and a weak acid (K 2 CO 3, Na 2 S - hydrolysis occurs at the anion)
SiO 3 2- + 2H 2 O H 2 SiO 3 + 2OH -
K 2 SiO 3 +2H 2 O H 2 SiO 3 +2KOH
The reaction of the medium is alkaline.
c) salts formed by a weak base and a weak acid (NH 4) 2 CO 3, Fe 2 (CO 3) 3 - hydrolysis occurs at the cation and at the anion.
2NH 4 + + CO 3 2- + 2H 2 O 2NH 3. H2O + H2CO3
(NH 4) 2 CO 3 + H 2 O 2NH 3. H2O + H2CO3
Often the reaction of the environment is neutral.
d) salts formed by a strong base and a strong acid (NaCl, Ba(NO 3) 2) are not subject to hydrolysis.
In some cases, hydrolysis proceeds irreversibly (as they say, it goes to the end). Thus, when mixing solutions of sodium carbonate and copper sulfate, a blue precipitate of hydrated basic salt precipitates, which, when heated, loses part of the water of crystallization and acquires green color– turns into anhydrous basic copper carbonate – malachite:
2CuSO 4 + 2Na 2 CO 3 + H 2 O (CuOH) 2 CO 3 + 2Na 2 SO 4 + CO 2
When mixing solutions of sodium sulfide and aluminum chloride, hydrolysis also proceeds to completion:
2AlCl 3 + 3Na 2 S + 6H 2 O 2Al(OH) 3 + 3H 2 S + 6NaCl
Therefore, Al 2 S 3 cannot be isolated from an aqueous solution. This salt is obtained from simple substances.
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. Similarly 
.
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.
|
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