The fourth period of the periodic system. General characteristics of d-elements D element 4 of the period of the periodic system corresponds to

DEFINITION

Potassium- the first element of the fourth period. It is located in group I of the main (A) subgroup of the Periodic Table.

Refers to the elements of the s - family. Metal. The metal elements included in this group are collectively called alkaline. Designation - K. Ordinal number - 19. Relative atomic mass - 39.102 a.m.u.

The electronic structure of the potassium atom

The potassium atom consists of a positively charged nucleus (+19), inside which there are 19 protons and 20 neutrons, and 19 electrons move around in 4 orbits.

Fig.1. Schematic structure of the potassium atom.

The distribution of electrons in orbitals is as follows:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 .

The outer energy level of the potassium atom contains 1 electron, which is a valence. The oxidation state of potassium is +1. The energy diagram of the ground state takes the following form:

Excited state despite the presence of vacant 3 p- and 3 d- there are no orbitals.

Examples of problem solving

EXAMPLE 1

The purpose of this work is to study the chemical properties of some transition metals and their compounds.

Metals of secondary subgroups, the so-called transition elements, belong to d-elements, since in their atoms d-orbitals are filled with electrons.

In transition metals, valence electrons are located in the d orbitals of the preexternal level and S orbitals of the outer electronic level. The metallicity of transition elements is explained by the presence of one or two electrons in the outer electron layer.

The incomplete d-sublevel of the preexternal electron layer determines the variety of valence states of the metals of the secondary subgroups, which in turn explains the existence of a large number of their compounds.

In chemical reactions, the d-orbital electrons participate after the S-electrons of the outer orbital are used. All or part of the electrons of the d orbitals of the penultimate electronic level can participate in the formation of chemical compounds. In this case, compounds corresponding to different valence states are formed. The variable valency of transition metals is their characteristic property (with the exception of metals of II and III secondary subgroups). Metals of subgroups IV, V, VI, VII of groups can be included in compounds both in the highest valence state (which corresponds to the group number) and in lower valence states. So, for example, titanium is characterized by 2-, 3-, 4-valence states, and manganese has 2-, 3-, 4-, 6- and 7-valence states.

Oxides and hydroxides of transition metals, in which the latter are in the lower valence state, usually exhibit basic properties, for example, Fe(OH) 2 . Higher oxides and hydroxides are characterized by amphoteric properties, for example TiO 2 , Ti(OH) 4 or acidic, for example
and
.

The redox properties of the compounds of the considered metals are also associated with the valence state of the metal. Combined with the lowest oxidation state usually exhibit reducing properties, and those with the highest oxidation state - oxidizing.

For example, for manganese oxides and hydroxides, the redox properties change as follows:

complex compounds.

A characteristic feature of transition metal compounds is the ability to form complexes, which is explained by the presence of a sufficient number of free orbitals in the metal ions in the outer and pre-outer electronic levels.

In the molecules of such compounds, a complexing agent is located in the center. Around it are coordinated ions, atoms or molecules called ligands. Their number depends on the properties of the complexing agent, the degree of its oxidation and is called the coordination number:

The complexing agent coordinates two types of ligands around itself: anionic and neutral. Complexes are formed when several different molecules are combined into one more complex one:

copper (II) sulfotetraamine; potassium hexacyanoferrate (III).

In aqueous solutions, complex compounds dissociate, forming complex ions:

The complex ions themselves are also capable of dissociation, but usually to a very small extent. For example:

This process proceeds reversibly and its equilibrium is sharply shifted to the left. Therefore, according to the law of mass action,

The constant Kn in such cases is called the instability constant of complex ions. The greater the value of the constant, the stronger the ability of the ion to dissociate into its constituent parts. The values ​​of Kn are given in the table:

Experience 1. Oxidation of Mn 2+ ions into ions
.

Put a little lead dioxide into the test tube so that only the bottom of the test tube is covered, add a few drops of concentrated
and one drop of solution
. Heat the solution and observe the appearance of ions
. Write an equation for the reaction. A solution of manganese salt should be taken in a small amount, since an excess of ions
restores
before
.

Experience 2. Oxidation with ions
in acidic, neutral and alkaline solutions.

Ion reduction products
are different and depend on the pH of the solution. So, in acidic solutions, the ion
reduced to ions
.

In neutral, slightly acidic and slightly alkaline solutions, i.e. in the pH range from 5 to 9, ion
is reduced with the formation of permanganic acid:

In strongly alkaline solutions and in the absence of a reducing agent, the ion
reduced to an ion
.

Pour 5-7 drops of potassium permanganate solution into three test tubes
. Add the same volume of dilute sulfuric acid to one of them, add nothing to the other, and a concentrated alkali solution to the third. In all three tubes, add dropwise, shaking the contents of the tube, a solution of potassium or sodium sulfite until the solution becomes colorless in the first tube, a brown precipitate forms in the second, and in the third the solution turns green. Write a reaction equation, keeping in mind that the ion
turns into ions
. Give an estimate of the oxidizing ability
in various environments according to the table of redox potentials.

Experience 3. Interaction of potassium permanganate with hydrogen peroxide. Place in a test tube 1 ml. hydrogen peroxide, add a few drops of sulfuric acid solution and a few drops of potassium permanganate solution. What gas is released? Test it with a smoldering torch. Write an equation for the reaction and explain it in terms of redox potentials.

Experience 4. Complex compounds of iron.

A) Obtaining Prussian blue. To 2-3 drops of an iron (III) salt solution, add a drop of acid, a few drops of water and a drop of a solution of hexatically - (P) potassium ferrate (yellow blood salt). Observe the appearance of a precipitate of Prussian blue. Write an equation for the reaction. This reaction is used to detect ions
. If a
take in excess, then instead of a precipitate of Prussian blue, its colloidal soluble form can form.

Investigate the relation of Prussian blue to the action of alkali. What is observed? which dissociates better. Fe (OH) 2 or complex ion
?

B) Obtaining iron thiocyanate III. To a few drops of iron salt solution, add a drop of potassium or ammonium thiocyanate solution
. Write an equation for the reaction.

Explore the ratio of thiocyanate
to alkalis and explain the observed phenomenon. This reaction, like the previous one, is used to detect the ion
.

Experience 5. Obtaining a complex compound of cobalt.

Place 2 drops of a saturated cobalt salt solution in a test tube and add 5-6 drops of a saturated ammonium solution: take into account that this forms a complex salt solution
. Complex ions
are colored blue, and hydrated ions
- in pink. Describe the observed phenomena:

1. Equation for obtaining a complex cobalt salt.

2. The equation of dissociation of the complex salt of cobalt.

3. Equation of dissociation of a complex ion.

4. Expression of the instability constant of the complex ion.

Control questions and tasks.

1. What properties (oxidizing or reducing) do compounds with the highest oxidation state of an element exhibit? Make an electron-ionic and molecular reaction equation:

2. What properties do compounds with an intermediate oxidation state of an element exhibit? Compose the electron-ionic and molecular reaction equations:

3. Specify the distinctive and similar properties of iron, cobalt, nickel. Why did D. I. Mendeleev place cobalt between iron and nickel in the periodic table of elements, despite the value of its atomic weight?

4. Write the formulas of complex compounds of iron, cobalt, nickel. What explains the good complexing ability of these elements?

5. How does the nature of manganese oxides change? What is the reason for this? What oxidation numbers can manganese have in compounds?

6. Are there similarities in the chemistry of manganese and chromium? What is it expressed in?

7. On what properties of manganese, iron, cobalt, nickel, chromium is their use in technology based?

8. Give an estimate of the oxidizing ability of ions
and reducing ability of ions
.

9. How to explain that the oxidation numbers of Cu, Ag, Au are greater than +17.

10. Explain the blackening of silver over time in air, the greening of copper in air.

11. Make an equation for the reactions proceeding according to the scheme.

d-elements and their compounds have a number of characteristic properties: variable oxidation states; ability to form complex ions; the formation of colored compounds.

Zinc is not among the transition elements. Its physical and chemical properties do not allow it to be classified as a transition metal. In particular, in its compounds it exhibits only one oxidation state and does not exhibit catalytic activity.

The d-elements have some peculiarities in comparison with the elements of the main subgroups.

1. In d-elements, only a small part of the valence electrons is delocalized throughout the crystal (whereas in alkali and alkaline earth metals, the valence electrons are completely given to collective use). The remaining d-electrons participate in the formation of directed covalent bonds between neighboring atoms. Thus, these elements in the crystalline state do not have a purely metallic bond, but a covalent metallic bond. Therefore, they are all solid (except for Hg) and refractory (with the exception of Zn, Cd) metals.

The most refractory metals are VB and VIB subgroups. In them, half of the d-sublevel is filled with electrons and the maximum possible number of unpaired electrons is realized, and, consequently, the largest number of covalent bonds. Further filling leads to a decrease in the number of covalent bonds and a drop in melting temperatures.

2. Due to the unoccupied d-shells and the presence of unoccupied ns- and np-levels close in energy, d-elements are prone to complex formation; their complex compounds are, as a rule, colored and paramagnetic.

3. d-Elements more often than the elements of the main subgroups, form compounds of variable composition (oxides, hydrides, carbides, silicides, nitrides, borides). In addition, they form alloys with each other and with other metals, as well as intermetallic compounds.

4. For d-elements, a large set of valence states is characteristic (Table 8.10) and, as a result of this, a change in acid-base and redox properties over a wide range.

Since some of the valence electrons are in s-orbitals, the lowest oxidation states they exhibit are usually equal to two. The exception is the elements whose ions E +3 and E + have stable configurations d 0 , d 5 and d 10: Sc 3+ , Fe 3+ , Cr + , Cu + , Ag + , Au + .

Compounds in which d-elements are in the lowest oxidation state form ion-type crystals, exhibit basic properties in chemical reactions and are, as a rule, reducing agents.

The stability of compounds in which d-elements are in the highest oxidation state (equal to the group number) increases within each transition row from left to right, reaching a maximum for 3d-elements for Mn, and in the second and third transition rows for Ru and Os, respectively. . Within one subgroup, the stability of compounds of the highest oxidation state decreases in the series 5d > 4d > 3d, as evidenced by the nature of the change in the Gibbs energy (isobaric-isothermal potential) of the same type of compounds, for example:

This phenomenon is due to the fact that with an increase in the principal quantum number within one subgroup, a decrease in the difference between the energies of the (n – 1)d- and ns-sublevels occurs. These compounds are characterized by covalent-polar bonds. They are acidic in nature and are oxidizing agents (CrO 3 and K 2 CrO 4 , Mn 2 O 7 and KMnO 4).

Compounds in which d-electrons are in intermediate oxidation states exhibit amphoteric properties and redox duality.

5. The similarity of d-elements with the elements of the main subgroups E(0) is fully manifested in the elements of the third group ns 2 np 1 and (n – 1)d 1 ns 2 . As the group number increases, it decreases; elements of subgroup VIIIA - gases, VIIIB - metals. In the first group, a distant similarity appears again (all elements are metals), and the elements of the IB subgroup are good conductors; this similarity is enhanced in the second group, since the d-elements Zn, Cd, and Hg do not participate in the formation of a chemical bond.

6. The d-elements of the IIIB–VIIB subgroups in higher oxidation states are similar in properties to the corresponding p-elements. Thus, in higher oxidation states, Mn (VII) and Cl (VII) are electronic analogues. The similarity of electronic configurations (s 2 p 6) leads to the similarity of the properties of compounds of heptavalent manganese and chlorine. Mn 2 O 7 and Cl 2 O 7 under normal conditions are unstable liquids, which are anhydrides of strong acids with the general formula HEO 4 . In the lower oxidation states, manganese and chlorine have a different electronic structure, which causes a sharp difference in the properties of their compounds. For example, lower chlorine oxide Cl 2 O (s 2 p 4) is a gaseous substance that is hypochlorous acid anhydride (HClO), while lower manganese oxide MnO (d 5) is a basic crystalline solid.

7. As you know, the reducing ability of a metal is determined not only by its ionization energy (M - ne - → M n +; + ∆H ionization), but also by the enthalpy of hydration of the formed cation (M n + + mH 2 O → M n + mH 2 O; –∆H hydr). The ionization energies of d-elements are high in comparison with other metals, but they are compensated by the large enthalpies of hydration of their ions. As a result, the electrode potentials of most d-elements are negative.

In the period with increasing Z, the reducing properties of metals decrease, reaching a minimum for group IB elements. Heavy metals of groups VIIIB and IB are called noble for their inertness.

The redox tendencies of compounds of d-elements are determined by the change in the stability of higher and lower oxidation states, depending on their position in the periodic system. Compounds with the maximum oxidation state of the element exhibit exclusively oxidizing properties, and with the lowest - reducing. Mn (OH) 2 is easily oxidized in air Mn (OH) 2 + 1 / 2O 2 \u003d MnO 2 + H 2 O. Mn (IV) compounds are easily reduced to Mn (II): MnO 2 + 4HCl \u003d MnCl 2 + Cl 2 + 2H 2 O, but strong oxidizing agents oxidize to Mn (VII). Permanganate ion MnO 4 - can only be an oxidizing agent.

Since for d-elements within the subgroup the stability of higher oxidation states increases from top to bottom, the oxidizing properties of compounds of the highest oxidation state fall sharply. So, chromium (VI) compounds (CrO 3, K 2 CrO 4, K 2 Cr 2 O 7) and manganese (VII) (Mn 2 O 7, KMnO 4) are strong oxidizing agents, and WO 3, Re 2 O 7 and salts of their respective acids (H 2 WO 4 , HReO 4) are difficult to recover.

8. The acid-base properties of d-element hydroxides are affected by the same factors (ionic radius and ion charge) as p-element hydroxides.

Hydroxides of lower oxidation states of d-elements usually exhibit basic properties, while those corresponding to higher oxidation states are acidic. In intermediate oxidation states, the hydroxides are amphoteric. The change in the acid-base properties of hydroxides with a change in the degree of oxidation is especially pronounced in manganese compounds. In the series Mn(OH) 2 - Mn(OH) 3 - Mn(OH) 4 - H 2 MnO 4 - HMnO 4, the properties of hydroxides change from the weak base Mn(OH) 2 through the amphoteric Mn(OH) 3 and Mn(OH) 4 to strong acids H 2 MnO 4 and HMnO 4 .

Within one subgroup, hydroxides of d-elements of the same oxidation state are characterized by an increase in basic properties when moving from top to bottom. For example, in the IIIB group, Sc (OH) 3 is a weak base, and La (OH) 3 is a strong base. Group IVB elements Ti, Zn, Hf form amphoteric hydroxides E(OH) 4 , but their acidic properties weaken when moving from Ti to Hf.

9. A distinctive feature of transition elements is the formation of phases of variable composition. These are, firstly, interstitial and substitutional solid solutions and, secondly, compounds of variable composition. Solid solutions are formed by elements with similar electronegativity, atomic radii, and identical crystal lattices. The more different elements are in nature, the less they dissolve in each other and the more prone to the formation of chemical compounds. Such compounds can have both constant and variable composition. Unlike solid solutions, in which the lattice of one of the components is preserved, compounds are characterized by the formation of a new lattice and new chemical bonds. In other words, only those phases of variable composition that sharply differ in structure and properties from the initial phases are classified as chemical compounds.

Compounds of variable composition are characterized by the following features:

a) The composition of these compounds depends on the method of preparation. So, depending on the synthesis conditions, titanium oxides have the composition TiO 1.2–1.5 and TiO 1.9–2.0; titanium and vanadium carbides - TiC 0.6–1.0 and VC 0.58–1.09, titanium nitride TiN 0.45–1.00.

b) Compounds retain their crystal lattice with significant fluctuations in the quantitative composition, that is, they have a wide area of ​​homogeneity. Thus, TiC 0.6–1.0, as follows from the formula, retains the titanium carbide lattice with a lack of up to 40% of carbon atoms in it.

c) The nature of the bond in such compounds is determined by the degree of filling of the d-orbitals of the metal. The electrons of the interstitial non-metal populate vacant d-orbitals, which leads to an increase in the covalence of the bonds. That is why the proportion of the metallic bond in the compounds of the initial elements of the d-series (groups IV–V) is reduced.

The presence of a covalent bond in them is confirmed by large positive enthalpies of formation of compounds, higher hardness and melting point, lower electrical conductivity compared to the metals that form them.

Copper is an element of the eleventh group of the fourth period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 29. It is designated by the symbol Cu (lat. Cuprum). The simple substance copper (CAS number: 7440-50-8) is a golden-pink ductile transition metal (pink in the absence of an oxide film). Since ancient times, it has been widely used by man.

Elements of the 4th Period of the Periodic Table

Rice. 3.8. Melting temperature dependence ( t pl) and boiling ( t bales), melting enthalpies (D H pl) and boiling (D H kip), Brinell hardness of simple substances of the 4th period on the number of electrons at the external energy level (the number of electrons in excess of a completely filled shell of the noble gas Ar)

As noted, to describe the chemical bond that occurs between metal atoms, one can use the representations of the method of valence bonds. The approach to the description can be illustrated by the example of a potassium crystal. The potassium atom has one electron in its outer energy level. In an isolated potassium atom, this electron is located at 4 s-orbitals. At the same time, in the potassium atom there are not very different in energy from 4 s-orbitals free, not occupied by electrons orbitals related to 3 d, 4p-sublevels. It can be assumed that during the formation of a chemical bond, the valence electron of each atom can be located not only on 4 s-orbitals, but also in one of the free orbitals. One valence electron of an atom allows it to realize one single bond with its nearest neighbor. The presence in the electronic structure of an atom of free orbitals that differ little in energy suggests that an atom can “capture” an electron from its neighbor to one of the free orbitals, and then it will be able to form two single bonds with its nearest neighbors. Due to the equality of distances to the nearest neighbors and the indistinguishability of atoms, various options for the implementation of chemical bonds between neighboring atoms are possible. If we consider a fragment of the crystal lattice of four neighboring atoms, then the possible options are shown in Fig. 3.9.

Elements of the 4th period of the Periodic Table - concept and types. Classification and features of the category "Elements of the 4th period of the Periodic Table" 2015, 2017-2018.

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