Patterns of changes in the electronegativity of elements in a group and period. Patterns of changes in electronegativity of elements in a group and period How electronegativity increases in a period

In this lesson you will learn about the patterns of changes in the electronegativity of elements in a group and period. Here you will look at what the electronegativity of chemical elements depends on. Using elements of the second period as an example, study the patterns of changes in the electronegativity of an element.

Topic: Chemical bond. Electrolytic dissociation

Lesson: Patterns of changes in the electronegativity of chemical elements in a group and period

Patterns of changes in relative electronegativity values ​​during the period

Let us consider, using the example of elements of the second period, the patterns of changes in the values ​​of their relative electronegativity. Fig.1.

Rice. 1. Patterns of changes in electronegativity values ​​of elements of period 2

The relative electronegativity of a chemical element depends on the charge of the nucleus and the radius of the atom. In the second period there are elements: Li, Be, B, C, N, O, F, Ne. From lithium to fluorine, the nuclear charge and the number of outer electrons increase. Number of electronic layers remains unchanged. This means that the force of attraction of outer electrons to the nucleus will increase, and the atom will seem to shrink. The atomic radius from lithium to fluorine will decrease. The smaller the radius of an atom, the stronger the outer electrons are attracted to the nucleus, which means the greater the value of relative electronegativity.

In the period with increasing nuclear charge, the radius of the atom decreases, and the value of relative electronegativity increases.

Rice. 2. Patterns of changes in the electronegativity values ​​of group VII-A elements.

Patterns of changes in relative electronegativity values ​​in the main subgroups

Let us consider the patterns of changes in the values ​​of relative electronegativity in the main subgroups using the example of elements of group VII-A. Fig.2. In the seventh group, the main subgroup contains halogens: F, Cl, Br, I, At. On the outer electron layer, these elements have the same number of electrons - 7. As the charge of the atomic nucleus increases during the transition from period to period, the number of electronic layers increases, and therefore the atomic radius increases. The smaller the atomic radius, the greater the electronegativity value.

In the main subgroup, with increasing charge of the atomic nucleus, the atomic radius increases, and the value of relative electronegativity decreases.

Since the chemical element fluorine is located in the upper right corner of D.I. Mendeleev’s Periodic Table, its relative electronegativity value will be maximum and numerically equal to 4.

Conclusion:Relative electronegativity increases with decreasing atomic radius.

In periods with increasing charge of the atomic nucleus, electronegativity increases.

In the main subgroups, as the charge of the atomic nucleus increases, the relative electronegativity of the chemical element decreases. The most electronegative chemical element is fluorine, as it is located in the upper right corner of D.I. Mendeleev’s Periodic Table.

Summing up the lesson

In this lesson, you learned about the patterns of changes in the electronegativity of elements in a group and period. On it you looked at what the electronegativity of chemical elements depends on. Using elements of the second period as an example, we studied the patterns of changes in the electronegativity of an element.

1. Rudzitis G.E. Inorganic and organic chemistry. 8th grade: textbook for general education institutions: basic level / G. E. Rudzitis, F.G. Feldman. M.: Enlightenment. 2011, 176 pp.: ill.

2. Popel P.P. Chemistry: 8th grade: textbook for general education institutions / P.P. Popel, L.S. Krivlya. -K.: IC “Academy”, 2008.-240 p.: ill.

3. Gabrielyan O.S. Chemistry. 9th grade. Textbook. Publisher: Bustard: 2001. 224s.

1. Nos. 1,2,5 (p. 145) Rudzitis G.E. Inorganic and organic chemistry. 8th grade: textbook for general education institutions: basic level / G. E. Rudzitis, F.G. Feldman. M.: Enlightenment. 2011, 176 pp.: ill.

2. Give examples of substances with a covalent nonpolar bond and an ionic bond. What is the significance of electronegativity in the formation of such compounds?

3. Arrange the elements of the second group of the main subgroup in order of increasing electronegativity.

Electronegativity (c) is the ability of an atom to hold external (type of atomic orbitals and the nature of their hybridization. valence) electrons. It is determined by the degree of attraction of these electrons to the positively charged nucleus.

This property manifests itself in chemical bonds as displacement of bond electrons towards a more electronegative atom.

The electronegativity of the atoms involved in the formation of a chemical bond is one of the main factors that determines not only the TYPE, but also the PROPERTIES of this bond, and thereby affects the nature of the interaction between atoms during a chemical reaction.

In L. Pauling's scale of relative electronegativities of elements (calculated based on the dependence of bond energies on differences in the electronegativities Dc of bonded atoms), metals and organogenic elements are arranged in the following row:

Electronegativity is not an absolute constant of an element. It depends on the effective charge of the atomic nucleus, which can change under the influence of neighboring atoms or groups of atoms

Element

The electronegativity of elements increases over the period and decreases somewhat in groups with increasing period number for elements of the I, II, V, VI and VII main subgroups, III, IV and V - secondary subgroups, has a complex dependence for elements of the III main subgroup (minimum EO in Al ), increases with increasing period number for elements of VII-VIII secondary subgroups. The s-elements of subgroup I have the lowest EO values, and the p-elements of groups VII and VI have the highest.

Examples of the electronic formula of p-elements

Aluminum Al 1s 2 2s 2 2p 6 3s 2 3p 1

Nitrogen N 1s 2 2s 2 2p 3

3(63) What is hybridization? consider the types of hydridization using the example of C2H6 molecules; s2n4; s2n2? Draw their electronic structure.

s-The orbital has a spherical shape, and R-orbital - the shape of a three-dimensional figure eight, oriented in a certain way in space.

When a carbon atom is excited, one of the two electrons 2 s-sublevel goes to free orbital 2 p-sublevel. This is possible due to the slight difference in energy 2 s- and 2 p-sublevels. Such a carbon atom in an excited state already has four unpaired electrons: one per 2 s- and three by 2 p-orbitals.

To justify the equivalence of the four valencies of the carbon atom, a description of its electronic structure is used using the concept of hybridization. This description is based on the idea that after mixing of orbitals, new, hybrid orbitals of equal energy are formed. In this case, hybridization should be understood as a mathematical, quantum-mechanical model, and not as a certain physical process.

For a carbon atom, three types of hybridization are possible (three valence states).

sp 3 -Hybridization- mixing one 2 s- and three 2 R-orbitals. All four hybrid orbitals are strictly oriented in space at an angle of 109°28" to each other, creating a geometric figure with thickened "petals" - a tetrahedron (Fig. 2). Therefore, sp3 The -hybridized carbon atom is often called "tetrahedral". State of the carbon atom with sp3-hybrid orbitals (first valence state) are characteristic of saturated hydrocarbons - alkanes.

sp 2 -Hybridization- mixing one 2s and two 2p-orbitals are not hybridized and are perpendicular to the plane in which the three sp2-hybrid orbitals. State of the carbon atom with sp2-hybrid orbitals (second valence state) are characteristic of unsaturated hydrocarbons of the ethylene series - alkenes.

sp -Hybridization- mixing of one 2s and one 2p orbital. Two hybrid orbitals are located in one straight line at an angle of 180° to each other. The remaining two non-hybridized 2p orbitals are located in mutually perpendicular planes. The state of the carbon atom with sp-hybrid orbitals (third valence state) is characteristic of unsaturated hydrocarbons of the acetylene series - alkynes.

The relationship between the type of orbital hybridization and the nature of the carbon atoms is shown in Table 1.

State of the carbon atom	Orbitals				Valence state of the carbon atom	Type of hybridization and molecular structure
Excited, initial state		2p x	2p y	2p z
Excited, in alkanes	2sp 3	2sp 3	2sp 3	2sp 3		(sp 3 ) Tetrahedral
Excited, in alkenes	2sp 2	2sp 2	2sp 2			(sp 2 ) Tetrahedral (planar)
Excited, in alkynes						(sp 2 ) Tetrahedral(linear)

An ethane molecule, for example, has seven σ bonds

located relative to each other at an angle of 109°28". The carbon atoms are in the first valence state (sp 3 hybridization).

In an ethylene molecule, as established using physical research methods, five σ bonds are located relative to each other at an angle of 120° and are in the same plane:

However, with this arrangement of bonds in ethylene, each carbon atom remains with one unpaired electron. They can no longer form a second σ bond between carbon atoms, since this would be accompanied by a violation of the Pauli principle. Therefore, such unpaired electrons of carbon atoms form a qualitatively different bond. The overlap of two electron clouds occurs in such a way that the eights of these clouds are perpendicular to the plane in which all six atoms of the ethylene molecule are located. This bond is called a π bond, and the electrons that form it are called π electrons. Since it is assumed that σ-bonds in an ethylene molecule are formed with the participation of hybridized electrons, and a π-bond with the participation of “pure” p-electrons (i.e., from the four electrons of each carbon atom, one s-electron and only two of the three p-electrons are hybridized), then the hybridization of electrons of a carbon atom in an ethylene molecule is denoted as sp 2 hybridization. In the ethylene molecule we encounter carbon atoms in the second valence state.

In the acetylene molecule, carbon is in the third valence state. In this molecule, all four atoms are located on the same straight line and the angles between σ bonds are 180° (sp hybridization). The electron clouds of two π-bonds are located along intersecting mutually perpendicular planes. Thus, according to the visual concepts presented above, two carbon-carbon bonds in a ethylene molecule and three in a ethylene molecule are not identical in their electronic structure. Until now, however, no chemical or physical experimental facts are known that could confirm this difference. The fact that carbon atoms connected by double bonds can easily add only two atoms of hydrogen, halogen, etc., so that one of the carbon-carbon bonds is preserved, is also easily consistent with the assumption that both carbon-carbon bonds are the same. In fact, if, for example, two hydrogen atoms are added along one of these identical carbon-carbon bonds, then as a result of this the nature of the second remaining carbon-carbon bond may change and differ from the two existing multiple bonds in the original compound, and may be stronger than each of them . Based on this assumption, it can be explained why a double bond easily adds only two, and a triple bond only four equivalents.

The inequality of bonds could be established by physical research methods, of which important is, for example, the determination of the energy of bonds C-C, C=C and C≡C. The energy values ​​of these bonds are respectively 79.3; 140.5; 196.7 kcal/mol

4.(93) Find the entropy, enthalpy and Gibbs energy under standard conditions for the following reaction: CO 2(g) →CO (G) + ½ O 2(g) and determine the possibility or impossibility of its occurrence under the same conditions.

Thermochemical calculations are based on Hess's law (1840): the thermal effect of the reaction depends only on the nature and physical state of the starting substances and final products, but does not depend on the transition path.

In thermochemical calculations they are used more often corollary of Hess's law : thermal effect of reaction (Δ H x.p) is equal to the sum of the enthalpies of formation Δ H arr of the reaction products minus the sum of the enthalpies of formation of the starting substances, taking into account the stoichiometric coefficients:

CO 2 (g) → CO (g) + ½ O 2 (g

Based on the table values ​​of ΔН 0 under standard conditions:

ΔН(СО 2(g))=-393.51 kJ/mol

ΔН(СО (g))= -110.53- (kJ/mol

ΔH(O 2(g))= 0

ΔH (c.r.) = -110.53-(-393.51) = 282.98 kJ/mol

Entropy is a function of state, i.e. its change (Δ S) depends only on the initial ( S 1) and final ( S 2) states and does not depend on the process path:

S((CO (g))= 197.55*10 -3 kJ/molK

S((CO 2(g))=213.66*10 -3 kJ/molK

S((O 2(g))=205.04*10 -3 kJ/molK

ΔS=(197.55*10 -3 +205.04*10 -3)- 213.66*10 -3 =188.93*10 -3 kJ/molK

Gibbs energy Δ G , can be found from the relation:

Δ G= Δ H– TΔ S.

ΔG=282.985-298*188.93*10 -3 =285.985-56.3=226.68>0

ΔG>0, therefore, under standard conditions the reaction is impossible.

5(123) How many times should the pressure be increased so that the rate of the reaction 2NO + O 2 → 2NO 2 increases 1000 times?

The dependence of the reaction rate on concentrations is determined law of mass action: at constant temperature the rate of chemical reaction is directly proportional to the product of the molar concentrations of the reactants.

V = K 2 [O 2 ]

An increase in pressure means a decrease in the volume of the gas mixture, and therefore an increase in the concentration of reactants

Let us denote the reaction rate before changing concentrations V 1, and after changing concentrations V 2

Let us denote the concentrations of substances by a and b:

then before the pressure increases –V 1 =k*a 2 b

after increasing pressure V 2 =k*(ax) 2 bx= k*a 2 bx3

From the conditions of the problem V 2 /V 1 =1000, we find x:

k*a 2 bx3/ k*a 2 b=1000

x 3 =1000, i.e. x=10

The pressure should be increased 10 times.

You can find out the activity of simple substances using the table of electronegativity of chemical elements. Denoted as χ. Read more about the concept of activity in our article.

What is electronegativity

The property of an atom of a chemical element to attract electrons from other atoms is called electronegativity. The concept was first introduced by Linus Pauling in the first half of the twentieth century.

All active simple substances can be divided into two groups according to physical and chemical properties:

• metals;
• non-metals.

All metals are reducing agents. In reactions they donate electrons and have a positive oxidation state. Nonmetals can exhibit reducing and oxidizing properties depending on their electronegativity value. The higher the electronegativity, the stronger the oxidizing properties.

Rice. 1. The actions of an oxidizing agent and a reducing agent in reactions.

Pauling created a scale of electronegativity. According to the Pauling scale, fluorine has the highest electronegativity (4), and francium the least (0.7). This means that fluorine is the strongest oxidizing agent and is able to attract electrons from most elements. On the contrary, francium, like other metals, is a reducing agent. It tends to give rather than accept electrons.

Electronegativity is one of the main factors that determines the type and properties of the chemical bond formed between atoms.

How to determine

The properties of elements to attract or give up electrons can be determined by the electronegativity series of chemical elements. According to the scale, elements with a value greater than two are oxidizing agents and exhibit the properties of a typical non-metal.

Item number	Element	Symbol	Electronegativity
	Strontium
	Ytterbium
	Praseodymium
	Prometheus
	Americium
	Gadolinium
	Dysprosium
	Plutonium
	Californium
	Einsteinium
	Mendelevium
	Zirconium
	Neptunium
	Protactinium
	Manganese
	Beryllium
	Aluminum
	Technetium
	Molybdenum
	Palladium
	Tungsten
	Oxygen

Substances with an electronegativity of two or less are reducing agents and exhibit metallic properties. Transition metals, which have variable oxidation states and belong to secondary subgroups of the periodic table, have electronegativity values ​​in the range of 1.5-2. Elements with an electronegativity equal to or less than one have pronounced reducing properties. These are typical metals.

In the electronegativity series, metallic and reducing properties increase from right to left, and oxidizing and nonmetallic properties increase from left to right.

Rice. 2. Electronegativity series.

In addition to the Pauling scale, you can find out how pronounced the oxidizing or reducing properties of an element are using the periodic table. Electronegativity increases in periods from left to right with increasing atomic number. In groups, the value of electronegativity decreases from top to bottom.

Rice. 3. Periodic table.

What have we learned?

Electronegativity shows the ability of an element to give or accept electrons. This characteristic helps to understand how pronounced the properties of an oxidizing agent (non-metal) or reducing agent (metal) are in a particular element. For convenience, Pauling developed an electronegativity scale. According to the scale, fluorine has the maximum oxidizing properties, and francium has the minimum. In the periodic table, the properties of metals increase from right to left and from top to bottom.

Test on the topic

Evaluation of the report

Average rating: 4.6. Total ratings received: 75.

Modern formulation of the Periodic Law, discovered by D. I. Mendeleev in 1869:

The properties of elements are periodically dependent on the ordinal number.

The periodically repeating nature of changes in the composition of the electronic shell of atoms of elements explains the periodic change in the properties of elements when moving through the periods and groups of the Periodic System.

Let us trace, for example, the change in higher and lower oxidation states of elements of groups IA – VIIA in the second – fourth periods according to Table. 3.

Positive All elements exhibit oxidation states except fluorine. Their values ​​increase with increasing nuclear charge and coincide with the number of electrons at the last energy level (with the exception of oxygen). These oxidation states are called highest oxidation states. For example, the highest oxidation state of phosphorus P is +V.

Negative oxidation states are exhibited by elements starting with carbon C, silicon Si and germanium Ge. Their values ​​are equal to the number of electrons missing up to eight. These oxidation states are called inferior oxidation states. For example, the phosphorus atom P at the last energy level is missing three electrons to eight, which means that the lowest oxidation state of phosphorus P is – III.

The values ​​of higher and lower oxidation states are repeated periodically, coinciding in groups; for example, in the IVA group, carbon C, silicon Si and germanium Ge have the highest oxidation state +IV, and the lowest oxidation state – IV.

This periodicity of changes in oxidation states is reflected in the periodic changes in the composition and properties of chemical compounds of elements.

A periodic change in the electronegativity of elements in the 1st-6th periods of groups IA–VIA can be similarly traced (Table 4).

In each period of the Periodic Table, the electronegativity of elements increases with increasing atomic number (from left to right).

In each group In the periodic table, electronegativity decreases as the atomic number increases (from top to bottom). Fluorine F has the highest, and cesium Cs has the lowest electronegativity among the elements of the 1st-6th periods.

Typical nonmetals have high electronegativity, while typical metals have low electronegativity.

Examples of tasks for parts A, B

1. In the 4th period the number of elements is equal to

2. Metallic properties of elements of the 3rd period from Na to Cl

1) get stronger

2) weaken

3) do not change

4) I don’t know

3. Nonmetallic properties of halogens with increasing atomic number

1) increase

2) decrease

3) remain unchanged

4) I don’t know

4. In the series of elements Zn – Hg – Co – Cd, one element not included in the group is

5. The metallic properties of elements increase in a number of ways

1) In – Ga – Al

2) K – Rb – Sr

3) Ge – Ga – Tl

4) Li – Be – Mg

6. Non-metallic properties in the series of elements Al – Si – C – N

1) increase

2) decrease

3) do not change

4) I don’t know

7. In the series of elements O – S – Se – Those sizes (radii) of an atom

1) decrease

2) increase

3) do not change

4) I don’t know

8. In the series of elements P – Si – Al – Mg, the dimensions (radii) of an atom are

1) decrease

2) increase

3) do not change

4) I don’t know

9. For phosphorus the element with less electronegativity is

10. A molecule in which the electron density is shifted towards the phosphorus atom is

11. Higher The oxidation state of elements is manifested in a set of oxides and fluorides

1) ClO 2, PCl 5, SeCl 4, SO 3

2) PCl, Al 2 O 3, KCl, CO

3) SeO 3, BCl 3, N 2 O 5, CaCl 2

4) AsCl 5, SeO 2, SCl 2, Cl 2 O 7

12. Lowest oxidation state of elements - in their hydrogen compounds and set fluorides

1) ClF 3, NH 3, NaH, OF 2

2) H 3 S + , NH +, SiH 4 , H 2 Se

3) CH 4, BF 4, H 3 O +, PF 3

4) PH 3, NF+, HF 2, CF 4

13. Valency for a multivalent atom is the same in a series of compounds

1) SiH 4 – AsH 3 – CF 4

2) PH 3 – BF 3 – ClF 3

3) AsF 3 – SiCl 4 – IF 7

4) H 2 O – BClg – NF 3

14. Indicate the correspondence between the formula of a substance or ion and the oxidation state of carbon in it

Electronegativity is the ability of atoms to displace electrons in their direction when forming a chemical bond. This concept was introduced by the American chemist L. Pauling (1932). Electronegativity characterizes the ability of an atom of a given element to attract a common electron pair in a molecule. Electronegativity values ​​determined by various methods differ from each other. In educational practice, they most often use relative rather than absolute values ​​of electronegativity. The most common is a scale in which the electronegativity of all elements is compared with the electronegativity of lithium, taken as one.

Among the elements of groups IA - VIIA:

electronegativity, as a rule, increases in periods (“from left to right”) with increasing atomic number, and decreases in groups (“from top to bottom”).

The patterns of changes in electronegativity among d-block elements are much more complex.

Elements with high electronegativity, the atoms of which have high electron affinity and high ionization energy, i.e., prone to the addition of an electron or the displacement of a pair of bonding electrons in their direction, are called nonmetals.

These include: hydrogen, carbon, nitrogen, phosphorus, oxygen, sulfur, selenium, fluorine, chlorine, bromine and iodine. According to a number of characteristics, a special group of noble gases (helium-radon) is also classified as nonmetals.

Metals include most of the elements of the Periodic Table.

Metals are characterized by low electronegativity, i.e., low ionization energy and electron affinity. Metal atoms either donate electrons to nonmetal atoms or mix pairs of bonding electrons from themselves. Metals have a characteristic luster, high electrical conductivity and good thermal conductivity. They are mostly durable and malleable.

This set of physical properties that distinguish metals from non-metals is explained by the special type of bond that exists in metals. All metals have a clearly defined crystal lattice. Along with atoms, its nodes contain metal cations, i.e. atoms that have lost their electrons. These electrons form a socialized electron cloud, the so-called electron gas. These electrons are in the force field of many nuclei. This bond is called metallic. The free migration of electrons throughout the volume of the crystal determines the special physical properties of metals.

Metals include all d and f elements. If from the Periodic Table you mentally select only blocks of s- and p-elements, i.e., elements of group A and draw a diagonal from the upper left corner to the lower right corner, then it turns out that non-metallic elements are located on the right side of this diagonal, and metallic ones - in the left. Adjacent to the diagonal are elements that cannot be unambiguously classified as either metals or non-metals. These intermediate elements include: boron, silicon, germanium, arsenic, antimony, selenium, polonium and astatine.

Ideas about covalent and ionic bonds played an important role in the development of ideas about the structure of matter, however, the creation of new physical and chemical methods for studying the fine structure of matter and their use showed that the phenomenon of chemical bonding is much more complex. It is currently believed that any heteroatomic bond is both covalent and ionic, but in different proportions. Thus, the concept of covalent and ionic components of a heteroatomic bond is introduced. The greater the difference in electronegativity of the bonding atoms, the greater the polarity of the bond. When the difference is more than two units, the ionic component is almost always predominant. Let's compare two oxides: sodium oxide Na 2 O and chlorine oxide (VII) Cl 2 O 7. In sodium oxide, the partial charge on the oxygen atom is -0.81, and in chlorine oxide -0.02. This effectively means that the Na-O bond is 81% ionic and 19% covalent. The ionic component of the Cl-O bond is only 2%.

List of used literature

1. Popkov V. A., Puzakov S. A. General chemistry: textbook. - M.: GEOTAR-Media, 2010. - 976 pp.: ISBN 978-5-9704-1570-2. [With. 35-37]
2. Volkov, A.I., Zharsky, I.M. Big chemical reference book / A.I. Volkov, I.M. Zharsky. - Mn.: Modern School, 2005. - 608 with ISBN 985-6751-04-7.