Periodic Transformation of Electron Shell Structure of Atoms of Elements

b. Group

 Elements s, p, d and f.

Elements whose final electron filling occurs in the s sub-shell are called s elements. Group IA and IIA elements are s elements.

The same definition holds for the p, d and f elements.

The elements in groups IIIA to VIIIA are p elements. The d elements are all in groups B.

The f elements have a special position: They can be grouped into groups IIIB, but most f elements have different properties from the group IIIB elements, so their properties are often examined separately.

Elements in which the final electron filling occurs at 4f are called lanthanides or lanthanide elements (with Z numbers from 58 to 71), and those in which the final electron filling occurs at 5f are called actinides or actinide elements (with Z numbers from 90 to 103).

The d and f elements are also known as the d and f transition elements.

 Group: Atoms of elements in the same group have similar valence electron configurations. This is the most basic factor that determines the similar properties of atoms, elements and compounds formed from elements in the same group.

The maximum oxidation number of most elements is equal to the group number (except fluorine, oxygen, group IB elements, most group VIIIB elements, lanthanides, actiodes, and noble gases). For example, the maximum oxidation number of group VA and group VB elements is +5.

 Group A: Atoms of group A elements have the following electron configuration characteristics:

- The final electron filling in atoms occurs in the s sublevel or the p sublevel.

For example: Atoms of element Z = 4: 1s 2 2s 2 belong to group A. Atoms of

Element Z = 31: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 all belong to group A.

- The number of electrons in the outermost layer of an atom is equal to the order number of the group containing it.

It.

For example: Atoms of element with Z = 31 belong to group IIIA.

To identify which group A an element belongs to, we rely on the electron configuration.

atoms as follows:

Group IA: The final electron filling of an atom ends at ns 1 (except

hydrogen

Group IIA: Final electron filling ends at ns 2 (except helium which has configuration

electron 1s 2 )

Group IIIA: Electron filling in atoms ends at np 1 Group IVA: Electron filling in atoms ends at np 2 Group VA: Electron filling in atoms ends at np 3 Group VIA: Electron filling in atoms ends at np 4 Group VIIA: Electron filling in atoms ends at np 5 Group VIIIA: Electron filling in atoms ends at np 6

This group has the additional element helium (Z = 2). The elements in group VIIIA are called noble gases.

 Group B: Group B elements have the following atomic electron configuration characteristics:

- The final electron filling in atoms of elements occurs in the d or f sub-level.

For example: Element Z = 30 has electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 , the last electron is filled in the 3d sublevel.

The element with Z = 59 has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10

5p 6 6s 2 4f 3 , the final electron filling in the 4f sublevel.

- The number of electrons in the outermost layer of group B atoms is less than 3 .

- The group number is equal to the sum of the number of electrons in the outermost layer and the number of electrons in the (n-1)d or (n-2)f sublayer (except for elements in groups IB, IIB, VIIIB) .

For example: Element with atomic number Z = 25 has electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6

4s 2 3d 5 is in group VIIB because it has 2 electrons in the 4th layer and 5 electrons in the 3d sub-layer.

To identify some elements in group B based on the following atomic electron configuration:

IIIB : Atoms of elements in this group have two outer electron sub-levels (n-1)d 1 ns 2 . People often group elements whose atoms are filled in (n-2)f into group IIIB. However, the properties of these elements are much different from those of group IIIB elements.

IVB : Atoms have two outermost sublevels: (n-1)d 2 ns 2 .

VB : Atoms have two outermost sub-levels (n-1)d 3 ns 2 (except niobium 4d 4 5s 1 ).

VIB : Atoms have two outermost sub-levels of (n-1)d 4 ns 2 (except Cr and Mo: (n-1)d 5

ns 1 )

Pt).

VIIB : Atoms have two outermost sublevels of (n-1)d 5 ns 2 .

VIIIB : Atoms have two outermost sub-levels of (n-1)d 6,7,8 ns 2 (except Ru, Rh, Pd,

IB : Atoms have two outermost sublevels of (n-1)d 10 ns 1 .

IIB : Atoms have two outermost sublevels of (n-1)d 10 ns 2 .

1.3.2. Periodic changes in the electron shell structure of atoms of elements

Comparing the electron shell structure of atoms of elements in different periods in the periodic table, we can draw the following observations:

Table 1.5: Changes in electron shell structure according to period

Period 1 : s1 s2

Period 2, 3: s 1 s 2 p 1 ……… p 6

Period 4, 5: s 1 s 2 d 1 ………………………………… d 10 p 1 ……… p 6 Period 6: s 1 s 2 d 1 , f 1 ……… … .f 14 , d 2 …………… d 10 p 1 ……… p 6 Period 7: s 1 s 2 d 1 , f 2 ………… ..f 14 ,d 2 , d 3 ………………………… .

- The period begins at the element whose first subshell of a new shell (s subshell) begins to have electrons and the period ends at the element whose p subshell of those shells is complete.

Because at the beginning of a period a new electron layer is formed, the period number that an element occupies is equal to the number of electron layers that an atom of that element has.

- Elements in the same subgroup (group A, group B) have similar electron shells (table 1.6) .

Table 1.6 . Variation in electron shell structure by group

Group A: IA

IIA	IIIA	IVA	VA	VIA	VIIA	VIIIA
s 1	s 2	s 2 p 1	s 2 p 2	s 2 p 3	s 2 p 4	s 2 p 5	s 2 p 6

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d 8 s 2

Group B: IIIB IVB VB VIB VIIB VIIIB IB IIB

d 1 s 2 d 2 s 2 d 3 s 2 d 5 s 1 d 5 s 2 d 6 s 2 d 7 s 2 d 10 s 1 d 10 s 2

f 1 s 2 …… .f 14 s 2

Lanthanide family , Actinide family

Thus, the number of electrons in the outermost layer of atoms of elements in general changes periodically as the nuclear charge increases.

We know that the electron structure in atoms of elements, especially the number of electrons in the outermost layer, determines the chemical properties of the elements. Therefore, the periodic variation of the number of electrons in the outermost layer determines the periodic properties of the elements and the compounds formed from those elements.

That is the content of Mendeleev's periodic law.

1.3.3. Periodic properties of atoms

a. Atomic and ionic radius (R)

The covalent atomic radius is half the distance between the nuclei of two identical atoms covalently bonded together at 25 0 C. For example, the distance between the two nuclei in the Cl2 molecule is 0.1998 nm (1nm = 10 -9 m), so the covalent atomic radius of chlorine is 0.0994 nm.

The atomic radius of a metal is half the distance between the nuclei of the two nearest metal atoms in a metal crystal. For example, the nearest distance between two sodium nuclei in a sodium crystal is 0.3716nm, so the atomic radius of sodium metal is 0.1858nm.

Ionic radius is calculated in ionic crystals . For example, the radius of the O 2- ion is

0.140nm and the radius of the F - ion is 0.136nm.

From left to right in a period, atomic radius generally decreases and in small periods atomic radius decreases faster than in large periods.

From top to bottom in a group A, the atomic and ionic radii increase , and in a group B from the first element to the second element these radii usually increase slowly, from the second element to the third element they usually do not change much.

b. Ionization energy of the atom (I)

Distinguish between the first ionization energy I 1 , the second ionization energy I 2 , the third ionization energy I 3 , …

The first ionization energy I 1 of an atom is the minimum energy required to remove an electron from an atom in the gaseous, elementary state into a 1+ ion also in the gaseous, elementary state:

Atom (k, cb) Ion + (k, cb) + e I 1 > 0

I is usually given in kJ/mol or eV (1eV is equivalent to 23.06 kcal/mol or 96.5 kJ/mol (Table 1.7).

For example: Ca (k, cb) Ca + (k, cb) + e I 1 = 590 kJ/mol

Table 1.7 . First ionization energy I 1 of some elements (kJ/mol)

Original

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Beige

I 1

1346

2372

520

899

800

1086

1043

1313

1680

2084

The second ionization energy I 2 corresponds to the process of removing the second electron from an ion carrying a charge of 1+ as follows:

Ca + (k,cb) Ca 2+ (k,cb) + e I 2 = 1145kJ/mol

Similar definition for third ionization energy (I 3 ), fourth ionization energy (I 4 ) … In an atom we always have: I 1 < I 2 < I 3 … < I n .

Ionization energy is a quantity that characterizes the ability of an atom to donate electrons when participating in an oxidation-reduction reaction.

From left to right in a period, the first ionization energy generally increases and reaches its maximum value at the last atom of the period (in noble gas atoms).

From the noble gas atom of the previous period to the first atom of the next period, the first ionization energy decreases abruptly, then increases gradually until the last atom of the period, similar to the previous period.

The process of variation of I 1 as above repeats from one cycle to another is called periodic variation of I 1 .

From top to bottom in group A, the value of I 1 decreases , while in group B this variation is slow and uneven, but usually decreases from top to bottom in a group.

c. Electron affinity (E)

Electron affinity is the energy released when an atom in the gaseous state, basically, gains one electron to become a negative ion in the gaseous state, basically, corresponding to the following process:

A(k, cb) + e A - (k, cb)

For example: Cl(k, cb) + e Cl - (k, cb) E 1 = - 348 KJ/mol The units of E are the same as those of I (kJ/mol) (Table 1.8).

Table 1.8. Electron affinity E of atoms of some elements (kJ/mol)

Original

Beige

-21

-18

123

-20

141

333

-55

Electron affinity indicates the oxidizing properties of an element. Electron affinity and ionization energy of an element vary in the same direction. As ionization energy increases, reducing properties decrease, oxidizing properties increase, so electron affinity increases.

In a period from left to right ionization energy and electron affinity increase.

From top to bottom, ionization energy and electron affinity decrease.

d. Electronegativity (X)

Electronegativity of an element is the ability of an atom of that element to attract electrons in a compound.

The element with high electronegativity will receive electrons from the element with low electronegativity.

than.

Electronegative elements have strong oxidizing properties, electronegative elements

small has strong reducing properties (properties of metals).

In principle, electronegativity has units of kJ/mol. However, relative electronegativity is used when comparing the electronegativity of an element with that of Li, so relative electronegativity has no units. Table 1.9 shows the electronegativity values of some elements.

Table 1.9 . Electronegativity of some elements