In the space around the nuclei in comparison with the distribution of electron density in the neutral atoms forming a given bond.
The so-called effective charges on atoms are used as a quantitative measure of bond polarity.
The effective charge is defined as the difference between the charge of electrons located in some region of space near the nucleus and the charge of the nucleus. However, this measure has only a conditional and approximate meaning, since it is impossible to unambiguously identify a region in a molecule that relates exclusively to an individual atom, and in the case of several bonds, to a specific bond.
The presence of an effective charge can be indicated by symbols of charges on atoms (for example, H + δ - Cl − δ, where δ is a certain fraction of the elementary charge).
Almost all chemical bonds, with the exception of bonds in diatomic homonuclear molecules, are polar to one degree or another. Covalent bonds are usually weakly polar. Ionic bonds are highly polar.
See also
Sources
Wikimedia Foundation. 2010.
- polar arrow
- Polar expeditions
See what “Polarity of chemical bonds” is in other dictionaries:
Polarity of chemical bonds- characteristics chemical bond(See Chemical bond), showing the redistribution of electron density in space near the nuclei in comparison with the initial distribution of this density in the neutral atoms forming this bond.... ...
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Valence- I Valence (from Latin valentia force) is the ability of an atom to form chemical bonds. A quantitative measure of strength is usually considered to be the number of other atoms in a molecule with which a given atom forms bonds. V. one of the fundamental... ... Great Soviet Encyclopedia
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Polarity.
Depending on the location of the common electron pair (electron density) between the nuclei of atoms, non-polar and polar bonds are distinguished.
A nonpolar bond is formed by atoms of elements with the same electronegativity. The electron density is distributed symmetrically relative to the atomic nuclei.
The bond between atoms with different electronegativity is called polar. The shared electron pair is shifted towards the more electronegative element. The centers of gravity of positive (b +) and negative (b -) charges do not coincide. The greater the difference in electronegativity of the elements forming a bond, the higher the polarity of the bond. When the electronegativity difference is less than 1.9, the bond is considered polar covalent.
For a diatomic molecule, the polarity of the molecule coincides with the polarity of the bond. In polyatomic molecules, the total dipole moment of a molecule is equal to the vector sum of the moments of all its bonds. The dipole vector is directed from + to –
Example 3. Using the valence bond method, determine the polarity of the tin(II) chloride and tin(IV) chloride molecules.
50 Sn belongs to p – elements.
Valence electrons 5s 2 5p 2. Distribution of electrons over quantum cells in the normal state:
| |
| |
| |
Chemical formulas tin (IV) chloride - SnCl 4, tin (II) chloride - SnCl 2
To construct the geometric shape of molecules, we depict the orbitals of unpaired valence electrons, taking into account their maximum overlap
Rice. 4. Geometric shape of SnCl 2 and SnCl 4 molecules
The electronegativity of Sn is 1.8. Cl – 3.0. Sn–Cl bond, polar, covalent. Let us depict the vectors of dipole moments of polar bonds.
in SnCl 2 and SnCl 4 molecules
SnCl 2 is a polar molecule
SnCl 4 is a non-polar molecule.
Substances, depending on temperature and pressure, can exist in gaseous, liquid and solid aggregate states.
In the gaseous state, substances are in the form of individual molecules.
In the liquid state in the form of aggregates, where the molecules are connected by intermolecular van der Waals forces or hydrogen bonds. Moreover, the more polar the molecules, the stronger the bond and, as a result, the higher the boiling point of the liquid.
IN solids structural particles are connected by both intramolecular and intermolecular bonds. Classify: ionic, metallic, atomic (covalent), molecular crystals and crystals with mixed bonds.
CONTROL TASKS
73. Why are the elements chlorine and potassium active, while the element argon, located between them, is low-active?
74. Using the method of valence bonds, explain why the water molecule (H 2 O) is polar, and the methane molecule (CH 4) is non-polar?
75. The substance carbon monoxide (II) is active substance, and carbon monoxide (IV) is classified as a low-active substance. Explain using the valence bond method.
76. How the strength of nitrogen and oxygen molecules changes. Explain using the valence bond method.
77. Why are the properties of a sodium chloride (NaCl) crystal different from those of a sodium (Na) crystal? What type of communication occurs in these crystals?
78. Using the valence bond method, determine the polarity of the molecules of aluminum chloride and hydrogen sulfide.
79. What type of hydroxides is rubidium hydroxide? Explain using the valence bond method.
80. The boiling point of liquid hydrogen fluoride is 19.5 0 C, and that of liquid hydrogen chloride (- 84.0 0 C). Why is there such a big difference in boiling points?
81. Using the method of valence bonds, explain why carbon tetrachloride (CCl 4) is nonpolar, and chloroform (CHCl 3) is a polar substance?
82. How does the strength of bonds change in CH 4 – SnH 4 molecules? Explain using the method of valence compounds.
83. What possible compounds form the elements: lead and bromine? Determine the polarity of these bonds.
84. Using the valence bond method, determine the polarity of nitrogen molecules and nitrogen (III) bromide.
85. The boiling point of water is 100 0 C, and that of hydrogen sulfide (60.7 0 C). Why is there such a big difference in boiling points?
86. Determine which compound has a stronger bond: tin bromide or carbon bromide? Determine the polarity of these connections.
87. Using the valence bond method, determine the polarity of the gallium iodide and bismuth iodide molecules.
88. Using the theory of chemical bonding, explain why xenon is a noble (low-active) element.
89. Indicate the type of hybridization (sp, sp 2, sp 3) in the compounds: BeCl 2, SiCl 4. Draw the geometric shapes of the molecules.
90. Draw the spatial arrangement of bonds in the molecules: boron hydride and phosphorus (III) hydride. Determine the polarity of molecules.
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In homonuclear molecules (H 2 , F 2 , etc.), the electron pair forming a bond belongs equally to each atom, therefore the centers of positive and negative charges in the molecule coincide. Such molecules are non-polar.
However, in heteronuclear molecules the contribution to the coupling of the wave functions of different atoms is different. Near one of the atoms, excess electron density appears, therefore, an excess negative charge, and near the other, a positive one. In this case, they talk about the displacement of an electron pair from one atom to another, but this should not be understood literally, but only as an increase in the probability of finding an electron pair near one of the nuclei of the molecule.
To determine the direction of such a shift and semi-quantitatively estimate its magnitude, the concept of electronegativity is introduced.
There are several scales of electronegativity. However, the elements are arranged in the same order in the electronegativity series, so the differences are insignificant, and the electronegativity scales are quite comparable.
Electronegativity according to R. Mulliken is half the sum of ionization energies and electron affinity (see section 2.10.3):
The valence electron pair shifts to the more electronegative atom.
It is more convenient to use relative rather than absolute values of electronegativity. The electronegativity of lithium 3 Li is taken as one. The relative electronegativity of any element A is equal to:
Heavy alkali metals have the lowest electronegativity (XFr = 0.7). The most electronegative element is fluorine (X F = 4.0). By periods, there is a general trend of increasing electronegativity, and by subgroups - its decrease (Table 3.4).
When using the data from this table in practice (as well as data from other electronegativity scales), it should be borne in mind that in molecules consisting of three or more atoms, the electronegativity value can change noticeably under the influence of neighboring atoms. Strictly speaking, a constant electronegativity cannot be assigned to an element at all. It depends on the valence state of the element, the type of compound, etc. Nevertheless, this concept is useful for a qualitative explanation of the properties of chemical bonds and compounds.
Table 3.4
Electronegativity of s- and p-elements according to Pauling
|
Period |
Group |
||||||
Bond polarity is determined by the displacement of the valence electron pair in diatomic molecules and is quantitatively characterized dipole moment, or electric dipole moment, molecules. It is equal to the product of the distance between the nuclei G in the molecule and the effective charge 5 corresponding to this distance:
Since G is considered a vector directed from a positive to a negative charge, the dipole moment is also a vector and has the same direction. The unit of dipole moment is debye D (1D = 3.33 Yu -30 C m).
The dipole moment of a complex molecule is defined as the vector sum of the dipole moments of all bonds. Therefore, if the AB molecule is symmetrical with respect to the line of each bond, the total dipole moment of such a molecule, despite the polarity
ness A-B connections, is equal to zero: D = ^ D; = 0. Examples include
live previously considered symmetrical molecules, the bonds in which are formed by hybrid orbitals: BeF 2, BF 3, CH 4, SF 6, etc.
Molecules in which bonds are formed by non-hybrid orbitals or hybrid orbitals involving lone pairs of electrons are asymmetrical with respect to the bond lines. The dipole moments of such molecules are not zero. Examples of such polar molecules: H 2 S, NH 3, H 2 0, etc. In Fig. Figure 3.18 shows a graphical interpretation of the summation of polar bond vectors in a symmetrical BeF 2 (fl) molecule and an asymmetrical H 2 S molecule (b).
Rice. 3.18. Dipole moments of BeF 2 (a) and H 2 S (b) molecules
As already noted, the greater the difference in electronegativity of the atoms forming a bond, the more strongly the valence electron pair shifts, the more polar the bond and, therefore, the greater the effective charge b, as illustrated in Table. 3.5.
Table 3.5
Change in the nature of the bond in a series of compounds of period II elements with fluorine
In a polar bond, two components can be roughly distinguished: ionic, due to electrostatic attraction, and covalent, due to overlapping orbitals. As the electronegativity difference increases OH the valence electron pair is increasingly shifted towards the fluorine atom, which acquires an increasingly negative effective charge. The contribution of the ionic component to the bond increases, and the share of the covalent component decreases. Quantitative changes become qualitative: in the UF molecule, the electron pair almost entirely belongs to fluorine, and its effective charge approaches unity, i.e. to the charge of the electron. We can assume that two ions were formed: the Li + cation and the anion F~, and the connection is due only to their electrostatic attraction (the covalent component can be neglected). This connection is called ionic. It can be considered as an extreme case of polar covalent bonding.
The electrostatic field has no preferred directions. That's why ionic bond unlike covalent directionality is not characteristic. An ion interacts with any number of ions of opposite charge. This accounts for another distinctive property of the ionic bond - lack of saturation.
For ionic molecules, binding energy can be calculated. If we consider ions as non-deformable balls with charges ±е, then the force of attraction between them depends on the distance between the centers of the ions G can be expressed by the Coulomb equation:
The energy of attraction is determined by the relation 
When approaching, a repulsive force appears due to the interaction of electron shells. It is inversely proportional to the distance to the power p:
Where IN- some constant. Exponent n significantly greater than unity and for various ion configurations lies in the range from 5 to 12. Taking into account that the force is the derivative of energy with respect to distance, from equation (3.6) we obtain:
With change G change Fnp And Fqtt. At some distance g 0 these forces are equalized, which corresponds to the minimum of the resulting interaction energy U Q . After the transformations you can get
This equation is known as the Born equation.
Minimum on the dependence curve U=f(r) correspond to the equilibrium distance r 0 and energy U Q . This is the binding energy between ions. Even if n is unknown, then we can estimate the binding energy by taking 1 /n equal to zero:
The error will not exceed 20%.
For ions with charges z l and z 2 equations (3.7) and (3.8) take the form:
Since in molecules of this type the existence of a bond approaching a purely ionic one is problematic, the last equations should be considered a very rough approximation.
At the same time, the problems of polarity and ionicity of bonds can be approached from the opposite position - from the point of view of polarization of ions. It is assumed that there is a complete transfer of electrons, and the molecule consists of isolated ions. Then the electron clouds shift under the influence electric field created by ions - polarization ions.
Polarization is a two-pronged process that combines polarizing effect ions from their polarizability. Polarizability is the ability of the electron cloud of an ion, molecule or atom to deform under the influence of the electrostatic field of another ion. The strength of this field determines the polarizing effect of the ion. From equation (3.10) it follows that the polarizing effect of an ion is greater, the greater its charge and the smaller its radius. The radii of cations are, as a rule, much smaller than the radii of anions, so in practice one often encounters the polarization of anions under the influence of cations, and not vice versa. The polarizability of ions also depends on their charge and radius. Ions of large size and charge are more easily polarized. The polarizing effect of an ion is reduced to drawing toward itself the electron cloud of an ion of opposite charge. As a result, the ionicity of the bond decreases, i.e. the bond becomes polar covalent. Thus, ion polarization reduces the degree of ionicity of the bond and has the opposite effect of bond polarization.
Polarization of ions in a molecule, i.e. An increase in the proportion of covalent bonds in it increases the strength of its decomposition into ions. In a series of compounds of a given cation with anions of the same type, the degree of dissociation in solutions decreases with increasing polarizability of the anions. For example, in the series of lead halides PbCl 2 - PbBr 2 - PI 2, the radius of the halide anions increases, their polarizability increases, and the decomposition into ions is weakened, which is reflected in a decrease in solubility.
When comparing the properties of salts with the same anion and sufficiently large cations, the polarization of the cations should be taken into account. For example, the radius of the Hg 2+ ion is larger than the radius of the Ca 2+ ion, so Hg 2+ is more polarized than Ca 2+. As a result, CaC1 2 is strong electrolyte, i.e. dissociates completely in solution, and HgCl 2 is a weak electrolyte, i.e. practically does not dissociate in solutions.
The polarization of ions in a molecule reduces its strength when it breaks down into atoms or molecules. For example, in the series CaCl 2 - CaBr 2 - Ca1 2, the radius of halide ions increases, their polarization by the Ca 2+ ion increases, and therefore the temperature of thermal dissociation into calcium and halogen decreases: CaHa1 2 = Ca + Ha1 2.
If an ion is easily polarized, then its excitation requires little energy, which corresponds to the absorption of visible light quanta. This is the reason for the coloration of solutions of such compounds. An increase in polarizability leads to an increase in color, for example, in the series NiCl 2 - NiBr 2 - Nil 2 (increased polarizability of the anion) or in the series KC1 - CuCl 2 (increased polarizability of the cation).
The boundary between polar covalent and ionic bonds is very arbitrary. For molecules in the gaseous state, it is believed that with a difference in electronegativity AH > 2,5 bond is ionic. In solutions of polar solvents, as well as in the crystalline state, the solvent molecules and neighboring particles at the nodes have a strong influence, respectively. crystal lattice. Therefore, the ionic nature of the bond appears at a significantly smaller difference in electronegativity. In practice, we can assume that the bond between typical metals and nonmetals in solutions and crystals is ionic.
On the hydrogen atom is +0.17, and on the chlorine atom -0.17.
The so-called effective charges on atoms are most often used as a quantitative measure of bond polarity.
The effective charge is defined as the difference between the charge of electrons located in some region of space near the nucleus and the charge of the nucleus. However, this measure has only a conditional and approximate [relative] meaning, since it is impossible to unambiguously identify a region in a molecule that relates exclusively to an individual atom, and in the case of several bonds, to a specific bond.
The presence of an effective charge can be indicated by symbols of charges on atoms (for example, H δ+ - Cl δ−, where δ is a certain fraction of the elementary charge) O − = C 2 + = O − (\displaystyle (\stackrel (-)(\mbox(O)))=(\stackrel (2+)(\mbox(C)))=(\stackrel (-)( \mbox(O))))(O δ− =C 2δ+ =O δ−), H δ+ -O 2δ− -H δ+ .
Almost all chemical bonds, with the exception of bonds in diatomic homonuclear molecules, are polar to one degree or another. Covalent bonds are usually weakly polar. Ionic bonds are highly polar.
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Subtitles
Effective charge
Values of relative effective charges obtained by various methods (optical spectroscopy, NMR, also based on quantum chemical calculations) may diverge. However, the available values of δ indicate that atoms in compounds of high charges do not have [corresponding to the absolute charge electron] and purely ionic compounds do not exist.
Instantaneous and induced dipoles.
A molecule is a dynamic system in which there is constant movement of electrons and vibration of nuclei. Therefore, the distribution of charges in it cannot be strictly constant. For example, the Cl 2 molecule is classified as non-polar: the value of its electric dipole moment is zero. However, at each given moment there is a temporary shift of charges to one of the chlorine atoms: Cl δ+ → Cl δ− or Cl δ− ← Cl δ+ with the formation instantaneous microdipoles. Since such a displacement of charges to any of the atoms is equally probable, the average distribution of charges exactly corresponds to the average zero value dipole moment.
For polar molecules, the value of the dipole moment at any given time is slightly greater or slightly less than its average value. The direction and magnitude of the instantaneous dipole are subject to continuous fluctuations in the permanent dipole moment. Thus, any non-polar and polar molecule (and an atom in it) can be considered as a set of periodic instantaneous microdipoles that change very quickly in magnitude and direction.
Rice. 32. Schemes of polar and non-polar molecules: a - polar molecule; b-nonpolar molecule
Every molecule contains both positively charged particles - atomic nuclei, and negatively charged ones - electrons. For each type of particle (or, more accurately, charge) one can find a point that will be, as it were, their “electric center of gravity.” These points are called the poles of the molecule. If the electrical centers of gravity of positive and negative charges in a molecule coincide, the molecule will be non-polar. Such, for example, are molecules H 2, N 2, formed by identical atoms that have common pairs of electrons in equally belong to both atoms, as well as many symmetrically constructed molecules with atomic bonds, for example methane CH 4, tetrachloride CCl 4.
But if the molecule is built asymmetrically, for example, it consists of two dissimilar atoms, as we have already said, the common pair of electrons can be shifted to a greater or lesser extent to the sideone of the atoms. It is obvious that in this case, due to the uneven distribution of positive and negative charges inside the molecule, their electrical centers of gravity will not coincide and the result will be a polar molecule (Fig. 32).
Polar molecules are
Polar molecules are dipoles. This term generally denotes any electrically neutral system, that is, a system consisting of positive and negative charges distributed in such a way that their electrical centers of gravity do not coincide.
The distance between the electrical centers of gravity of those and other charges (between the poles of the dipole) is called the length of the dipole. The length of the dipole characterizes the degree of polarity of the molecule. It is clear that the dipole length is different for different polar molecules; The larger it is, the more pronounced the polarity of the molecule is.
Rice. 33. Schemes of the structure of CO2 and CS2 molecules
In practice, the degree of polarity of certain molecules is determined by measuring the so-called dipole moment of the molecule m, which is defined as the product of the length of the dipole l to the charge of its pole e:
t =l e
The magnitudes of dipole moments are associated with certain properties of substances and can be determined experimentally. Order of magnitude T always 10 -18, since the electric charge
throne is equal to 4.80 10 -10 electrostatic units, and the length of the dipole is a value of the same order as the diameter of the molecule, i.e. 10 -8 cm. Below are the dipole moments of the molecules of some inorganic substances.
Dipole moments of some substances
T 10 18
. . . .. …….. 0
Water……. 1.85
. . . ………..0
Hydrogen chloride……. 1.04
Carbon dioxide…….0
Bromide. …… 0.79
Carbon disulfide…………0
Hydrogen iodide…….. 0.38
Hydrogen sulfide………..1.1
Carbon monoxide……. 0,11
Sulfur dioxide. . . ……1.6
Hydrocyanic acid……..2.1
Determining the values of dipole moments allows one to draw many interesting conclusions regarding the structure of various molecules. Let's look at some of these findings.
Rice. 34. Scheme of the structure of a water molecule
As one would expect, the dipole moments of hydrogen and nitrogen molecules are zero; the molecules of these substances are completelysymmetrical and therefore electric charges they are evenly distributed. The lack of polarity in carbon dioxide and carbon disulfide shows that their molecules are also built symmetrically. The structure of the molecules of these substances is shown schematically in Fig. 33.
Somewhat unexpected is the presence of a rather large dipole moment near water. Since the formula of water is similar to the formulas of carbon dioxide
and carbon disulfide, one would expect that its molecules would be built in the same waysymmetrically, like the molecules CS 2 and CO 2.
However, in view of the experimentally established polarity of water molecules (polarity of molecules), this assumption must be discarded. Currently, the water molecule is credited with an asymmetrical structure (Fig. 34): two hydrogen atoms are connected to an oxygen atom so that their bonds form an angle of about 105°. A similar arrangement of atomic nuclei is found in other molecules of the same type (H 2 S, SO 2) that have dipole moments.
The polarity of water molecules explains many of its physical properties.