We are preparing for the exam in chemistry. Oxygen and hydrogen compounds of nonmetals. Brief description of their properties The acidic properties of hydrogen compounds increase in the series

Acidic properties are those that are most pronounced in a given environment. There are a whole range of them. It is necessary to be able to determine the acidic properties of alcohols and other compounds not only to identify the content of the corresponding environment in them. This is also important for recognizing the substance being studied.

There are many tests for acidic properties. The most elementary is immersion in the substance of an indicator - litmus paper, which reacts to the hydrogen content by turning pink or red. Moreover, a more saturated color demonstrates a stronger acid. And vice versa.

Acid properties increase with increasing radii of negative ions and, consequently, of the atom. This ensures easier removal of hydrogen particles. This quality is characteristic feature strong acids.

There are the most characteristic acidic properties. These include:

Dissociation (elimination of a hydrogen cation);

Decomposition (formation of water under the influence of temperature and oxygen);

Interaction with hydroxides (resulting in the formation of water and salt);

Interaction with oxides (as a result, salt and water are also formed);

Interaction with metals preceding hydrogen in the activity series (salt and water are formed, sometimes with the release of gas);

Interaction with salts (only if the acid is stronger than the one that formed the salt).

Chemists often have to produce their own acids. There are two ways to remove them. One of them is mixing acid oxide with water. This method is used most often. And the second is the interaction of a strong acid with a salt of a weaker one. It is used somewhat less frequently.

It is known that acidic properties are manifested in many. They can be more or less expressed depending on K. The properties of alcohols are manifested in the ability to abstract a hydrogen cation when interacting with alkalis and metals.

Alcoholates - salts of alcohols - are capable of hydrolyzing under the influence of water and releasing alcohol with metal hydroxide. This proves that the acidic properties of these substances are weaker than those of water. Consequently, the environment is expressed more strongly in them.

The acidic properties of phenol are much stronger due to the increased polarity of the OH compound. Therefore, this substance can also react with hydroxides of alkaline earth and alkali metals. As a result, salts are formed - phenolates. To identify phenol, it is most effective to use with (III), in which the substance acquires a blue-violet color.

So, acidic properties in different compounds manifest themselves in the same way, but with different intensities, which depends on the structure of the nuclei and the polarity of hydrogen bonds. They help determine the environment of a substance and its composition. Along with these properties, there are also basic ones, which increase with the weakening of the first.

All these characteristics appear in most complex substances and form an important part of the world around us. After all, it is at their expense that many processes take place not only in nature, but also in living organisms. Therefore, acidic properties are extremely important; without them, life on earth would be impossible.

Modern wording periodic law : properties simple substances, as well as the forms and properties of compounds of elements are periodically dependent on the magnitude of the charge of the nuclei of their atoms (ordinal number).

Periodic properties are, for example, atomic radius, ionization energy, electron affinity, electronegativity of the atom, as well as some physical properties elements and compounds (melting and boiling points, electrical conductivity, etc.).

The expression of the Periodic Law is

periodic table elements .

The most common version of the short form of the periodic table, in which the elements are divided into 7 periods and 8 groups.

Currently, the nuclei of atoms of elements up to number 118 have been obtained. The name of the element with serial number 104 is rutherfordium (Rf), 105 – dubnium (Db), 106 – seaborgium (Sg), 107 – bohrium (Bh), 108 – hassium (Hs ), 109 – meitnerium ( Mt), 110 - darmstadtium (Ds), 111 - roentgenium (Rg), 112 - copernicium (Cn).
On October 24, 2012, in Moscow, at the Central House of Scientists of the Russian Academy of Sciences, a solemn ceremony was held to assign the name “flerovium” (Fl) to the 114th element, and “livermorium” (Lv) to the 116th element.

Periods 1, 2, 3, 4, 5, 6 contain 2, 8, 8, 18, 18, 32 elements, respectively. The seventh period is not completed. Periods 1, 2 and 3 are called small, the rest - big.

In periods from left to right, metallic properties gradually weaken and non-metallic properties increase, since with an increase in the positive charge of atomic nuclei, the number of electrons in the outer electronic layer increases and a decrease in atomic radii is observed.

At the bottom of the table are 14 lanthanides and 14 actinides. IN lately Lanthanum and actinium began to be classified as lanthanides and actinides, respectively.

Groups are divided into subgroups - main ones, or subgroups A and side effects, or subgroup B. Subgroup VIII B – special, it contains triads elements that make up the families of iron (Fe, Co, Ni) and platinum metals (Ru, Rh, Pd, Os, Ir, Pt).

From top to bottom in the main subgroups, metallic properties increase and non-metallic properties weaken.

The group number usually indicates the number of electrons that can participate in the formation chemical bonds. This is the physical meaning of the group number. Elements of side subgroups have valence electrons not only in their outer layers, but also in their penultimate layers. This is the main difference in the properties of the elements of the main and secondary subgroups.

Periodic table and electronic formulas of atoms

To predict and explain the properties of elements, you must be able to write electronic formula atom.

In an atom located in ground condition, each electron occupies a vacant orbital with the lowest energy. The energy state is determined primarily by temperature. The temperature on the surface of our planet is such that the atoms are in the ground state. At high temperatures, other states of atoms, which are called excited.

The sequence of arrangement of energy levels in order of increasing energy is known from the results of solving the Schrödinger equation:

1s< 2s < 2p < 3s < Зр < 4s 3d < 4p < 5s 4d < 5p < 6s 5d 4f < 6p.

Let's consider the electronic configurations of atoms of some elements of the fourth period (Fig. 6.1).



Rice. 6.1. Distribution of electrons over the orbitals of some elements of the fourth period

It should be noted that there are some features in electronic structure atoms of elements of the fourth period: for atoms Cr and C u by 4 s-shell contains not two electrons, but one, i.e. "failure" external s -electron to the previous one d-shell.

Electronic formulas of 24 Cr and 29 Cu atoms can be represented as follows:

24 Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1,

29 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 .

The physical reason for the “violation” of the filling order is associated with the different penetrating ability of electrons into the inner layers, as well as the special stability of the electronic configurations d 5 and d 10, f 7 and f 14.

All elements are divided into four types

:

1. In atoms s-elements filled in s - outer layer shell ns . These are the first two elements of each period.

2. In atoms p-elements electrons fill the p-shells of the outer np level . These include the last 6 elements of each period (except the first and seventh).

3. U d-elements filled with electrons d -sublevel of the second outside level ( n-1)d . These are elements of intercalary decades of large periods located between s- and p-elements.

4. U f-elements filled with electrons f -sublevel of the third outside level ( n-2)f . These are lanthanides and actinides.

Changes in the acid-base properties of element compounds by groups and periods of the periodic system
(Kossel diagram)

To explain the nature of the change in the acid-base properties of compounds of elements, Kossel (Germany, 1923) proposed using a simple scheme based on the assumption that there is a purely ionic bond in the molecules and a Coulomb interaction takes place between the ions. The Kossel scheme describes the acid-base properties of compounds containing E–H and E–O–H bonds, depending on the charge of the nucleus and the radius of the element forming them.

Kossel diagram for two metal hydroxides (for LiOH and KOH molecules ) is shown in Fig. 6.2. As can be seen from the presented diagram, the radius of the Li ion + less than the ion radius K+ and OH The - - group is bonded more tightly to the lithium ion than to the potassium ion. As a result, KOH will be easier to dissociate in solution and the basic properties of potassium hydroxide will be more pronounced.



Rice. 6.2. Kossel diagram for LiOH and KOH molecules

In a similar way, you can analyze the Kossel scheme for two bases CuOH and Cu(OH) 2 . Since the radius of the Cu ion 2+ less, and the charge is greater than that of the ion Cu+, OH - - the group will be held more firmly by the Cu 2+ ion .
As a result, the base Cu(OH)2 will be weaker than CuOH.

Thus, the strength of bases increases as the radius of the cation increases and its positive charge decreases .

Kossel diagram for two oxygen-free acids HCl and HI shown in Fig. 6.3.



Rice. 6.3. Kossel diagram for HCl and HI molecules

Since the radius of the chloride ion is smaller than that of the iodide ion, the H+ ion more strongly bound to the anion in the hydrochloric acid molecule, which will be weaker than hydroiodic acid. Thus, the strength of anoxic acids increases with increasing radius of the negative ion.

The strength of oxygen-containing acids changes in the opposite way. It increases as the radius of the ion decreases and its positive charge increases. In Fig. Figure 6.4 shows the Kossel diagram for two acids HClO and HClO 4.



Rice. 6.4. Kossel diagram for HClO and HClO 4

Ion C1 7+ is firmly bound to the oxygen ion, so the proton will be more easily split off in the HC1O molecule 4 . At the same time, the bond of the C1 ion+ with O 2- ion less strong, and in the HC1O molecule the proton will be more strongly retained by the O anion 2- . As a result, HClO 4 is a stronger acid than HClO.

Thus, An increase in the oxidation state of an element and a decrease in the radius of the element’s ion increase the acidic nature of the substance. On the contrary, a decrease in the oxidation state and an increase in the ion radius enhance the basic properties of substances.

Examples of problem solving

Compose electronic formulas of the zirconium atom and ions O 2– , Al 3+ , Zn 2+ . Determine what type of elements the Zr, O, Zn, Al atoms belong to.

40 Zr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 2 5s 2,

O 2– 1s 2 2s 2 2p 6,

Zn 2+ 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 ,

Al 3+ 1s 2 2s 2 2p 6 ,

Zr – d-element, O – p-element, Zn – d-element, Al – p-element.

Arrange the atoms of the elements in order of increasing their ionization energy: K, Mg, Be, Ca. Justify the answer.

Solution. Ionization energy– the energy required to remove an electron from an atom in the ground state. In the period from left to right, the ionization energy increases with increasing nuclear charge; in the main subgroups from top to bottom it decreases as the distance from the electron to the nucleus increases.

Thus, the ionization energy of atoms of these elements increases in the series K, Ca, Mg, Be.

Arrange atoms and ions in increasing order of their radii: Ca 2+, Ar, Cl –, K +, S 2– . Justify the answer.

Solution. For ions containing the same number of electrons (isoelectronic ions), the radius of the ion will increase as its positive charge decreases and its negative charge increases. Consequently, the radius increases in the order Ca 2+, K +, Ar, Cl –, S 2–.

Determine how the radii of ions and atoms change in the series Li + , Na + , K + , Rb + , Cs + and Na, Mg, Al, Si, P, S.

Solution. In the series Li + , Na + , K + , Rb + , Cs + the radius of ions increases as the number of electronic layers of ions of the same sign with a similar electronic structure increases.

In the series Na, Mg, Al, Si, P, S, the radius of atoms decreases, since with the same number of electron layers in the atoms, the charge of the nucleus increases, and, therefore, the attraction of electrons by the nucleus increases.

Compare the strength of acids H 2 SO 3 and H 2 SeO 3 and bases Fe(OH) 2 and Fe(OH) 3.

Solution. According to the Kossel scheme H 2 SO 3 more strong acid, than H 2 SeO 3 , since the ion radius SE 4+ greater than the ion radius S 4+, which means the S 4+ – O 2– bond is stronger than bond Se 4+ – O 2– .

According to the Kossel scheme Fe(OH)

2 stronger base since the radius of the Fe ion 2+ more than Fe ion 3+ . In addition, the charge of the Fe ion 3+ greater than that of Fe ion 2+ . As a result, the Fe bond 3+ – О 2– is stronger than Fe 2+ – O 2– and ION – easier to split off in a molecule Fe(OH)2.

Problems to solve independently

6.1.Compose electronic formulas for elements with a nuclear charge of +19, +47, +33 and those in the ground state. Indicate what type of elements they belong to. What oxidation states are characteristic of an element with a nuclear charge of +33?




6.2.Write the electronic formula of the Cl ion – .

General properties of main classes inorganic compounds. Conditions for the occurrence of “exchange reactions”.

1. Acid-base properties of hydrogen compounds.

A) Comment on the ability of water to self-ionize (equation, K W). Based on the structure of molecules (their polarizability), explain the patterns of changes in solubility in water and acid-base properties of the corresponding solutions of methane (CH 4), ammonia (NH 3), hydrogen fluoride (HF) and hydrogen chloride (HCl). Make up the necessary equations.

b) Using the concept of the polarizing effect of cations on the H–O bond, and also taking into account the number of hydroxo groups, explain the pattern of changes in the acid-base properties of hydroxides LiOH–Be(OH) 2 –H 3 BO 3 –H 2 CO 3 –HNO 3 –H 3 PO 4 –H 2 SO 4 –(H 2 SeO 4)–HClO 4. Make up dissociation equations for the proposed substances.

2. Mandatory and optional(including special ones) reactions of acids and bases.

A) Which of the following substances (solutions) can 20% solutions of nitric, sulfuric and acetic acids react with: solutions of KOH, NH 3, H2S; Zn(OH)2, H3PO2; BaCl 2 and crystalline Cu, Ca 3 (PO 4) 2 .

b) Which of the following substances (solutions) can 20% solutions of potassium hydroxide and ammonia react with: solutions of H 2 SO 4, CH 3 COOH; Zn(OH)2, Al(OH) 3 ; MgCl 2 and crystalline Ag2O, AgCl.

In both versions of the experiment, the formulas of substances are highlighted in bold, the interaction with which will require writing non-obvious equations.

The task involves only a theoretical discussion, but... The reaction equations must be thought out and written in advance, including in ionic form.

3. Conditions for exchange reactions with salts.

What exchange reactions can be performed using the proposed reagents: dilute solutions MnSO4, Ba(NO3)2, saturated solution SrSO 4, crystalline CuS And FeS, as well as concentrated solutions of HCl, CO 2 and NH 3. Consider the possibility of performing reactions that require the participation of salt. Justify your proposals by calculating the constants of the corresponding exchange equilibria. Consider possible signs of reactions occurring.

It must be borne in mind that if substances that are sparingly soluble in water are used as a reagent (in in this case CuS and FeS), then reactions with their participation must necessarily be accompanied by dissolution, i.e. the products of such reactions should not themselves produce precipitation. For example, it is illiterate to think through the reaction of FeS ↓ and H 2 CO 3 in the hope of obtaining a FeCO 3 precipitate.

Reactions with rich solution SrSO 4 suggest the use of solution over the precipitate, and not the sediment itself.

4. Dependence of pH of solutions on the composition of salts.

Determine the hydrolyzability of the ions of the proposed salts (NH 4 NO 3, KCl, CH 3 COONa, Na 2 CO 3, AlCl 3, CH 3 COONH 4),

· create equations for the hydrolysis of an ion (ions, if both the cation and the anion of the salt are involved in the hydrolysis); calculate the hydrolysis constant ( TO G (Al 3+) take equal to ~10 -5).

write an equation in molecular form

(make a molecular equation based on the predominant ionic reaction ).

· Arrange the salts in order of increasing hydrolyzability.

Test hydrolyzability experimentally. To do this, pour ~1 ml of the corresponding solution into a clean test tube, moisten a glass rod in this solution and apply the solution to indicator paper. Use the color scale to estimate the approximate pH value of the solution. Why in two cases does the pH correspond to a neutral environment?

5. Medium in solutions of medium and acidic salts.

Write down the equations of the predominant ionic reactions, affecting the environment in solutions of potassium ortho-, hydro- and dihydrogen phosphate (K 3 PO 4, K 2 HPO 4, KN 2 PO 4). It must be borne in mind that in solutions of acidic salts, in addition to hydrolysis reactions, dissociation of the anions H 2 PO 4 ‒ and HPO 4 2 ‒ also takes place. The environment will be determined by the predominant reaction. Compare the constants of the competing reactions of hydrolysis and dissociation of anions and draw a conclusion about the pH (more or less than 7). Compare the results of the preliminary analysis with the actual pH value (determine using a universal indicator).

Reference data for preparing for experiments 3, 4, 5

3. Periodic law and periodic table chemical elements

3.4. Periodic changes in the properties of substances

The following properties of simple and complex substances change periodically:

• the structure of simple substances (initially non-molecular, for example from Li to C, and then molecular: N 2 - Ne);
• melting and boiling temperatures of simple substances: when moving from left to right along the period, t pl and t bp initially, in general, increase (diamond is the most refractory substance), and then decrease, which is associated with a change in the structure of simple substances (see above);
• metallic and non-metallic properties of simple substances. Over the period, with increasing Z, the ability of atoms to give up an electron decreases (E and increases), accordingly, the metallic properties of simple substances weaken (non-metallic properties increase, since E avg of atoms increases). From top to bottom in groups A, on the contrary, the metallic properties of simple substances increase, and non-metallic properties weaken;
• composition and acid-base properties of oxides and hydroxides (Table 3.1–3.2).

Table 3.1

Composition of higher oxides and simplest hydrogen compounds of A-group elements

As can be seen from table. 3.1, the composition of higher oxides changes smoothly in accordance with the gradual increase in covalency (oxidation state) of the atom.

As the charge of the atomic nucleus increases in a period, the basic properties of oxides and hydroxides weaken, and the acidic properties increase. The transition from basic oxides and hydroxides to acidic ones in each period occurs gradually, through amphoteric oxides and hydroxides. As an example in table. Figure 3.2 shows the change in the properties of oxides and hydroxides of elements of the 3rd period.

Table 3.2

Oxides and hydroxides formed by elements of the 3rd period and their classification

In groups A, as the charge of the atomic nucleus increases, the basic properties of oxides and hydroxides increase. For example, for group IIA we have:

1. BeO, Be(OH) 2 - amphoteric (weak basic and acidic properties).

2. MgO, Mg(OH) 2 - weak, basic properties.

3. CaO, Ca(OH) 2 - pronounced basic properties (alkalis).

4. SrO, Sr(OH) 2 - pronounced basic properties (alkalies).

5. BaO, Ba(OH) 2 - pronounced basic properties (alkalis).

6. RaO, Ra(OH) 2 - pronounced basic properties (alkalies).

The same trends can be traced for elements of other groups (for the composition and acid-base properties of binary hydrogen compounds, see Table 3.1). In general, with increasing atomic number over the period, the basic properties of hydrogen compounds weaken, and the acidic properties of their solutions increase: sodium hydride dissolves in water to form an alkali:

NaH + H 2 O = NaOH + H 2,

and aqueous solutions of H 2 S and HCl are acids, with hydrochloric acid being the stronger.

1. In groups A, as the charge of the atomic nucleus increases, the strength of oxygen-free acids also increases.

2. In hydrogen compounds, the number of hydrogen atoms in a molecule (or formula unit) first increases from 1 to 4 (groups IA–IVA), and then decreases from 4 to 1 (groups IVA–VIIA).

3. Volatile (gaseous) at ambient conditions. are only hydrogen compounds of elements of groups IVA–VIIA (except H 2 O and HF)

The described trends in changes in the properties of atoms of chemical elements and their compounds are summarized in table. 3.3

Table 3.3

Changes in the properties of atoms of elements and their compounds with increasing charge of the atomic nucleus

Example 3.3. Specify the formula of the oxide with the most pronounced acidic properties:

Solution. The acidic properties of oxides increase from left to right across the period, and weaken from top to bottom across group A. Taking this into account, we come to the conclusion that the acidic properties are most pronounced in the oxide Cl 2 O 7.

Answer: 4).

Example 3.4. The element anion E 2− has the electronic configuration of an argon atom. Specify the formula of the highest oxide of an element's atom:

Solution. The electronic configuration of the argon atom is 1s 2 2s 2 2p 6 3s 2 3p 6, therefore the electronic configuration of the E atom (the E atom contains 2 electrons less than the E 2− ion) is 1s 2 2s 2 2p 6 3s 2 3p 4, which corresponds to the atom sulfur. The element sulfur is in the VIA group, the formula of the highest oxide of elements of this group is EO 3.

Answer: 1).

Example 3.5. Indicate the symbol of the element whose atom has three electron layers and forms a volatile (no.) compound of the composition EN 2 (H 2 E):

Solution. Hydrogen compounds of the composition EN 2 (H 2 E) form atoms of elements of groups IIA and VIA, but are volatile at zero conditions. are compounds of group VIA elements, which include sulfur.

Answer: 3).

The characterized trends in changes in the acid-base properties of oxides and hydroxides can be understood based on the analysis of the following simplified diagrams of the structure of oxides and hydroxides (Fig. 3.1).

From a simplified reaction scheme

it follows that the efficiency of the interaction of the oxide with water to form a base increases (according to Coulomb’s law) with increasing charge on the E n + ion. The magnitude of this charge increases as the metallic properties of the elements increase, i.e. from right to left across the period and from top to bottom across the group. It is in this order that the basic properties of the elements increase.

Rice. 3.1. Scheme of the structure of oxides (a) and hydroxides (b)

Let us consider the reasons underlying the described changes in the acid-base properties of hydroxides.

With an increase in the oxidation state of the element +n and a decrease in the radius of the E n + ion (this is precisely what is observed with an increase in the charge of the nucleus of an element’s atom from left to right across the period), the E–O bond is strengthened, and the O–H bond is weakened; the process of hydroxide dissociation according to the acid type becomes more probable.

From top to bottom in the group, the radius E n + increases, but the value n + does not change, as a result, the strength of the E–O bond decreases, its breaking becomes easier, and the process of dissociation of the hydroxide according to the main type becomes more likely.