Examples of strong and weak acids and bases. Strong electrolytes include acid. A strong electrolyte is co2 o2 h2s h2so4

How to distinguish strong electrolytes from weak ones? and got the best answer

Answer from Pavel Beskrovny[master]
STRONG ELECTROLYTES, when dissolved in water, almost completely dissociate into ions. For such electrolytes, the VALUE OF THE DEGREE OF DISSOCIATION tends to UNITY in dilute solutions.
Strong electrolytes include:
1) almost all salts;
2) strong acids, for example: H2SO4 (sulfuric acid), HCl (hydrochloric acid), HNO3 (nitric acid);
3) all alkalis, for example: NaOH (sodium hydroxide), KOH (potassium hydroxide).
WEAK ELECTROLYTES, when dissolved in water, almost do not dissociate into ions. For such electrolytes, the VALUE OF THE DEGREE OF DISSOCIATION tends to ZERO.
Weak electrolytes include:
1) weak acids - H2S (hydrogen sulfide), H2CO3 (carbonic acid), HNO2;
2) aqueous solution of ammonia NH3 * H2O
DEGREE OF DISSOCIATION is the ratio of the number of particles disintegrated into ions (Nd) to the total number of dissolved particles (Np) (denoted by the Greek letter alpha):
a= Nd / Nr. Electrolytic dissociation is a reversible process for weak electrolytes. I hope you know what electrolytes are, since you’re asking. This is simpler, if it’s more complicated, then see above (for a number of EOs).
Electrolytic dissociation is a reversible process for weak electrolytes.
If you have questions, then go to soap.

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

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

Strong electrolytes include:

Soluble salts;

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

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

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

Therefore, in the dissociation equations of weak electrolytes, the reversibility sign (⇄) is used.

Weak electrolytes include:

Almost all organic acids and water;

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

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

Ionic reaction equations

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

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

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

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

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

If strong electrolytes are not involved in the reaction, the ionic form of the equation is absent:

Mg(OH)2(s) + 2HF(r) = MgF2↓ + 2H2O

TICKET No. 23

Hydrolysis of salts

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

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

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

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

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

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

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

Question 24 Classification of oxides

Oxides are called complex substances whose molecules include oxygen atoms in oxidation state - 2 and some other element.

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

Salt-forming oxides For example,

CuO + 2HCl → CuCl 2 + H 2 O.

CuO + SO 3 → CuSO 4.

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

CuO + 2HCl → CuCl 2 + H 2 O.

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

CuO + SO 3 → CuSO 4.

Non-salt-forming oxides These are oxides that do not form salts. Examples include CO, N 2 O, NO.

ELECTROLYTES– substances whose solutions or melts conduct electric current.

NON-ELECTROLYTES– substances whose solutions or melts do not conduct electric current.

Dissociation– decomposition of compounds into ions.

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

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

When writing equations for the dissociation of strong electrolytes, an equal sign is used.

Strong electrolytes include:

· Soluble salts ( see solubility table);

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

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

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

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

Weak electrolytes include:

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

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

· Insoluble metal hydroxides (Mg(OH) 2 , Fe(OH) 2 , Zn(OH) 2) ( look at the grounds-cweak electrolytes in the solubility table).

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

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

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

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

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

H3PO4⇄H++H2PO−4 (first stage, α 1),

H2PO−4⇄H++HPO2−4 (second stage, α 2),

НPO2−4⇄Н++PO3−4 (third stage, α 3).

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

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

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

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

HCOOH ⇄ HCOO − + H +

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

HCOONa → HCOO−+Na+

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

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

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

where is the degree of dissociation.

Ostwald's dilution law was derived by W. Ostwald in 1888 and he also confirmed it experimentally. The experimental establishment of the correctness of Ostwald's dilution law was of great importance for substantiating the theory of electrolytic dissociation.

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

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)