The origin and nature of chemical bonds


Химия и фармакология

The bond dissociation energy, also called simplythe bond energy, is a measure of its strength. The bond dissociation energy is always positive because otherwise the chemical bond would spontaneously break with the liberation of energy. A condition for the formation of a chemical bond is a decrease in the potential energy of a system of interacting atoms.



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The origin and nature of chemical bonds

When atoms interact, achemical bond may appear between them that leads to the formation of a stable polyatomic system—a molecule, molecular ion, or crystal.

The electrostatic force  of attraction which hold atoms, ions and molecules together, are, in general,  referred to aschemical bonds.

Since nuclei of atoms bear positive  charges, we can say that all chemical bonds are formed as the result of simultaneous attraction of two or more nuclei for electrons.

The bond dissociation energy, also called simplythe bond energy, is a measure of its strength. The bond dissociation energy is always positive because otherwise the chemical bond would spontaneously break with the liberation of energy.A condition for the formation of a chemical bond  is a decrease in the potential energy of a system of interacting atoms.

Basic types of chemical bonds. Covalent bond and ionic bond.

The electrostatic force of attraction between the oppositely charged ions is calledionic bonds.The force of attraction between ions may be expressed by the inverse square law

F =

where k is constant of proportionality, q1 and q2 are the charges of the ions and r is interatomic distance. A large  number of substances classified as ionic salts exists. They conduct electric current in the liquid phase but not in the solid phase. Ionic crystals have relatively high melting  and boiling points, are brittle and are easily broken when a stress is exerted on them.

In 1916, an American scientist Gilbert Lewis (1875-1946) suggested that atoms in molecules might be held together by an electrostatic force attraction between atomic nuclei andelectron-pairs shared by the bonding atoms. Lewis called an electron-pair bond between atoms acovalent bond. He proposed that chemical bonds form between electrons residing in the outermost, orvalence, shell of each bonding atom. Thevalence electronsare the electrons in the outermost shell on an atom that can be gained or lost in a chemical reaction. We shall consider a covalent bond in detail below.

Type of chemical bonds depends ondifference in electronegativity  between atoms.Ionic-bond model explain the formation of positive and negative ions in terms of electron transfer from an atom having a low electronegativity to one having a high electronegativity. If the difference in of Pauling electronegativity values between the two atoms is greater than 1.7, then the bond is predominately ionic; if the difference is less than 1.7, then it is predominately covalent; if it is 1.7, then the bond has approximately 50% ionic and 50% covalent character. We can say  about a partial ionic character of covalent bondin moleculeswith increasing of the difference in electronegativity  between atoms because the more electronegative atoms acquires a partial negative charge and the less electronegative acquires a partial positive charge.

The bonding between atoms in metals is different from these two general types and is called ametallic bond.The outer electrons from each atom are free to move from one atom to another through the lattice. The term 'delocalized' is often used to describe their condition. A metal lattice can therefore be viewed as an array of cations held together by a cloud of delocalized outer-shell electrons.

A metallic bondis the electrostatic force of attraction that two neighbouring cations have for delocalized electrons between them

Hydrogen bond is a specific bond of the lightest element hydrogen with active non-metals such as nitrogen, oxygen, and fluorine.


There are several definitions of valence expressing the different aspects of this concept. Originally the valence of a hydrogen atom was adopted as the unit of valence.

The valence of another element can be expressed by the number of hydrogen atoms which one atom of this other element combines with or substitutes.

Thus, in the compounds HC1, H2O, NH3, and CH4, the hydrogen referred valence of chlorine is one, of oxygen is two, of nitrogen is three, and of carbon is four.

The valence can be expressed in other ways, for example,  by the number of chemical bonds formed by an atom of the given element  or by the number of atoms directly surrounding a given atom.

Most elements display a different valence in different compounds. For instance, carbon forms two oxides with oxygen: carbon monoxide CO and carbon dioxide CO2. The valence of carbon in the monoxide is two, and in the dioxide four. As a rule, the valence of an element cannot be characterised by a single number.

Donor-acceptor way of forming a covalent bond

A covalent bond in which one atom donates both bonding electrons is called coordinate covalent bond

Let us consider, for example, the electron structure of an ammonia molecule:

Here the dots stand for electrons originally belonging to the nitrogen atom, and the crosses stand for the electrons originally belonging to the hydrogen atoms. Of the eight outer electrons of the nitrogen atom, six form three covalent bonds and are common for the nitrogen atom and the hydrogen atoms. But two electrons belong only to the nitrogen and form an unshared electron pair. Such a pair of electrons can also participate in the formation of a covalent bond with another atom if the outer electron layer of this atom has a free orbital, for example, in the hydrogen ion:

Therefore, when a molecule of NH3 interacts with a hydrogen ion, a covalent bond is produced between them. The unshared pair of electrons of the nitrogen atom becomes common for the two atoms, as a result of which anammonium ion NH4 is formed:

Here, the covalent bond was produced at the expense of the pair of electrons originally belonging to one atom (to the donor of the electron pair), and of the free orbital of another atom (the acceptorof the electron pair). This is the so-called donor-acceptor way of forming a covalent bond. In the example considered, the nitrogen atom is the donor of the electron pair, and the hydrogen atom is its acceptor.

It has been established experimentally that the four N—H bonds in the ammonium ion are equivalent in all respects.

The total number of covalent bonds which a given atom can form is restricted. It is determined bythe total number of valence orbitals.Quantum-mechanical calculations show that such orbitals include the s and p orbitals of the outer electron layer and the d orbitals of the preceding layer.

Covalent Bond. The Method of Valence Bonds

We know that a quantum-mechanical description of the structure of molecules is obtained, as for many-electron atoms, only on the basis of approximate solutions of the Schrodinger equation.Such an approximate calculation was performed for the first timefor a hydrogen moleculeby W. HeitlerandF. London in 1927. They first considered asystem of two hydrogen atoms at a great distance from each other.

For this condition, it is possible to take into account only the interaction of each electron with its "own" nucleus and to express the dependence of the wave function of the system being considered on the coordinates.

Heitler and London assumed further that the dependence of the wave function on the coordinates which they had found is also retained when the hydrogen atoms approach each other. Here, however, it is already necessary to take into consideration also the interactions (between the nuclei, between the electrons, etc.) that could be disregarded when the atoms are far apart.

As a result, Heitler and London obtained equations allowing them to find the dependence of the potential energyE of a system consisting of two hydrogen atoms on the distance r between the nuclei of these atoms. The results of the calculations depend on whether the spins of the interacting electrons are identical (parallel) or opposite (antiparallel).When the spins are identical (Fig. 1, curve a), approaching of the atoms leads to a continuous growth in the energy of the systemandno chemical bond appears between the atoms.

Whenthe spins are opposite (Fig. 1, curve b),approaching of the atoms to a certain distance ro is attended by diminishing of the system's energy.When r=ro, the system has the lowest potential energy, that is  in the most stable state. A molecule of H2 is formed with opposite spins of the atom electrons.

The formation of a chemical bond between the hydrogen atoms is the result of mutual penetration (“overlapping”) of the electron clouds when the interacting atoms approach each other (Fig. 2). This attraction  predominates over the mutual repulsion of the electrons with a like charge, and the result is the formation of a stable molecule.

Thus, the chemical bond in a hydrogen molecule is achieved by the formation of a pair of electrons with opposite spins belonging to both atoms. Such a two-electron bond between two centres is called a covalent bond (or, sometimes, an electron pair bond).

The concepts of the mechanism of chemical bond using the example of the hydrogen molecule were extended to more intricate molecules. The theory of the chemical bond developed on this basis was named the method of valence bonds (the VB method).

The VB method is based on the following propositions:

A covalent chemical bond is formed by two electrons with anti-parallel spins, this electron pair belonging to two atoms.

Combinations of such two-electron-bonds between two centres reflecting the electron structure of a molecule are known as electron dot formulas (or Lewis electron formulas).

The strength of a covalent bond grows with an increasing degree of overlapping of the interacting electron clouds.

Some possible variants of electron cloud overlapping with indication of the signs of the relevant wave functions are depicted inFig. 3. Positive overlapping (Fig. 3, a)occurs when the absolute value of the wave function for the net electron cloud formed will be greater than the values ofj for the isolated atoms. Negative overlapping resulting in the mutual repulsion of the nuclei (Fig. 3b).

Using the electron dot (Lewis) structures we can represent the formation of a hydrogen molecule as follows:

H×  +×H® H: H

This formula shows that when two hydrogen atoms combine into a molecule, each of the atoms acquires a stable two-electron shell like the electron shell of the helium atom.

Similar formulas can be used to show the formation of a nitrogen molecule:

The structure of the molecules of some compounds such as ammonia, water, carbon dioxide, and methane can be depicted by the formulas:

Non-polar and polar covalent bond

A covalent bond is called non-polar or homopolar if electron pair is distributed symmetrically in space relative to the both atoms (the moleculesH2,N2, andCl2). A covalent bond is called polar or heteropolar if a diatomic molecule consists of atoms of different elements, the common electron cloud is displaced towards one of the atoms. When a covalent bond is formed between two atoms of different elements, the common electron cloud is displaced to the atom having a higher electronegativity. The more electronegative atoms acquires a partial negative charge and the less electronegative acquires a partial positive charge. These charges are calledeffective chargesof the atoms in  a molecule.

A polar covalent bond in molecule, for example, in ahydrogen chloride molecule  can be considered as aelectric dipoles. Although the net charge of a dipole is zero, an electric field is formed in the space surrounding it as shown inFig. 4. The strength of this field is proportional to the dipole moment,m, of the molecule, which is the product of the absolute value of the chargeq of an electron and the distancel between the centres of the positive and negative charges in the molecule:

m= ql

The dipole moment of a molecule is a quantitative measure of its polarity. The dipole moments of molecules are usually measured in debyes (D): 1 D = 3.33× 10-30 C.m.

A molecule is more polar, the more the shared electron pair is displaced towards one of the atoms, i.e. the higher are the effective charges of the atoms and the greater is the length of the dipole. For instance, the dipole moments ofHC1, HBr, and HI are respectively 1.04, 0.79, and 0.38 D, which is associated with the reduction in the difference between the electronegativities of the atoms when transferring from HCl to HBr and HI.

Polyatomic molecules may also benon-polar—with a symmetrical distribution of the charges, orpolar—with an asymmetrical distribution of the charges. Here account must be taken not only of the magnitude of the dipole moment, but also of itsdirection, i.e. the dipole moment of each bond should be treated as a vector. Hence, the net dipole moment of the molecule equal to the vector sum of the dipole moments of the individual bonds. A dipole moment is usually considered to be directed from the positive end of the dipole to its negative one.

Figure 4 shows possible shapes of a type AB3 molecule. If the molecule is constructed in the form of a plane triangle (Fig. 4a), the vector sum of the dipole moments of the individual bonds is zero. If the molecule has a pyramidal structure (Fig. 4b), its net dipole moment differs from zero—the molecule is polar. Thus, the moleculeBF3 whose dipole moment is zero has a plane structure, whereas the polar molecule NH3 (m == 1.46 D) is constructed in the form of a pyramid.

The polarity of molecules appreciably affects the properties of the substances they form. Polar molecules tend to orient themselves relative to one another with oppositely charged ends. A consequence of suchdipole-dipole interaction is the mutual attraction of polar molecules and strengthening of the bonds between them. This is whysubstances formed of polar molecules have, as a rule, higher melting and boiling points than substances whose molecules are non-polar.

Ion-dipole interaction  between polar molecules or having an ionic structure substance and polar liquid promotes the electrolytic dissociation in solution formed.

Peculiarities of a covalent bond: direction  and saturability

The ability of atoms to participate in the formation of a restricted number of covalent bonds has been called the saturability of covalent bonding.

The overlapping of the valence electron clouds of the interacting atoms in a covalent bond formation  is possible only with a definite mutual orientation of the electron clouds. In other words, a covalent bond isdirectional.

In a hydrogen molecule(see Fig. 2 ), the atomic s-electron clouds overlap  near the bond axis. A covalent bond formed in such a way is known as as-bond.If p-electron clouds are oriented along the bond axis, they can also participate in the formation of a sigma bond as in the molecule F2(Fig. 5). Thus, in the molecule HF(Fig. 5), a covalent sigma bond is produced owing to overlapping of the 1s-electron cloud of the hydrogen atom and the 2p-electron cloud of the fluorine atom.

When p-electron clouds oriented perpendicularly to the bond axis interact, two overlap regions are formed at both sides of this axis instead of a single region. Such a covalent bond is called apbond.

Let us consider the formation of the nitrogen molecule N2 (Fig. 6). Each nitrogen atom has three unpaired2p electronswhose electron clouds are oriented in three mutually perpendicular directions.

Nitrogen atoms are bound in the molecule N2 by three covalent bonds. But these bonds are not equivalent: one of them is a sigma bond, and the other two arepbonds.

The concept of covalent bonds being directional makes it possible to explain the mutual arrangement of the atoms in polyatomic molecules. Thus, in the formation ofa water molecule the p-electron clouds of the oxygen atom are oriented in mutually perpendicular directions, the molecule H2O has, as shown in a V-shaped structure, and we can expect the angle between the 0—H bonds to be 90 deg.

The moleculeNH3 formed upon the interaction of the threep electrons of a nitrogen atom with thes electrons of three hydrogen atoms  has the shape of a pyramid with the nitrogen atom at its apex and the hydrogen atoms at the apices of its base.In this case too, we can expect the angles between the N—H bonds to be 90 deg.

These conclusions on the mutual arrangement of the atoms in the molecules NH3 and H2O correspond to actual facts. But the angles between the bonds, however (the valence angles) differ from 90 deg: in the water molecule the angle HOH is 104.3 deg, and in the ammonia molecule the angle HNH is 107.8 deg.

To explain why the valence angles in the molecules H2O and NH3 differ from 90 deg, we must take into account  the notion of atomic orbital hybridization.

Hybridization of atomic electron orbitals

The method of hybridization of atomic orbitals proceeds from the assumption thatin the formation of a molecule, instead of the initial atomic s-, p-, and d-electron clouds,equivalent "blended" or hybridelectron clouds are formed that are stretched out in a direction towards the neighbouring atoms, the result being their more complete overlapping with the electron clouds of these atoms. More complete overlapping of the valence electron clouds leads to the formation of a stronger chemical bond and, consequently, to an additional gain in energy. If this gain in energy is sufficient to more than compensate the expenditure of energy for the deformation of the initial atomic electron clouds.

Let us consider as an example of hybridization the formation of a molecule of beryllium fluoride BeF2. Each fluorine atom in this molecule has one unpaired electron that participates in the formation of a covalent bond. A beryllium atom in the unexcited state (1s22s2) has no unpaired electrons. Therefore, to participate in the formation of chemical bonds, the beryllium atom must pass over into the excited state(1s22s12p1):

The excited atom Be* that is formed has two unpaired electrons:

When s- and p- electron clouds of Be* atom overlap with the p-electron clouds of the two fluorine atoms, covalent bonds may form.

As already mentioned above, two equivalent hybrid orbitals(sp hybrids) may be formed instead of the initials andp orbitals of the beryllium atom. The shape and arrangement of these orbitals are shown inFig. 8.

The extended shape of the hybrid orbitals results in more complete overlapping of the interacting electron clouds, and, consequently, stronger chemical bonds are formed. The above example of hybridization is calledsp hybridization.Thesp orbitals are oriented in opposite directions, which results in a linear structure of the molecule and both Be—F bonds are equivalent in all respects.

For example, in the hybridization of ones and twop orbitals(sp2 hybridization—read "ess-pea-two"), three equivalentsp2 orbitals are formed. In this case, the hybrid electron clouds are arranged in directions that are in one plane and oriented at 120 deg to one another. It is evident that the formation of a planar triangular molecule corresponds to this type of hybridization.

The boron trifluoride molecule BF3 is an example of a molecule withsp2 hybridization. All three B—F bonds in the molecule BF3 are equivalent.

If ones and threep orbitals participate in hybridization(sp3 hybridization), the result is the formation of foursps hybrid orbitals extended towards the corners of a tetrahedron, i.e. oriented at 109 deg 28 min relative to one another (Fig. 9). Such hybridization occurs, for instance, in an excited carbon atom


when a methane molecule CH4 is being formed. This is why the methane molecule has the shape of a tetrahedron, all four C—H bonds in this molecule being equivalent.

Let us return to the consideration of the water molecule structure. Its formation is attended bysp3 hybridization of the oxygen atomic orbitals. It is exactly for this reason that the valence angle HOH in the H2O molecule (104.3 deg) is close not to 90 deg, but to the tetrahedral angle (109.5 deg). The slight deviation of this angle from 109.5 deg can be understood if we take into account the non-equivalence of the state of the electron clouds. Indeed, in a methane molecule all eight electrons occupying thesp3 hybrid orbitals in the carbon atom participate in the formation of the covalent bonds C—H.

In a water molecule (II), only four of the eight electrons occupying thesp3 hybrid orbitals of the oxygen atom form 0—H bonds, while two electron pairs remain unshared, i.e. belong only to the oxygen atom. This leads to a certain asymmetry in the distribution of the electron clouds surrounding the oxygen atom and, consequently, to deviation of the angle between the 0—H bonds from 109.5 deg. In the formation of an ammonia molecule,sp3 hybridization of the atomic orbitals of the central atom (nitrogen) also occurs. It is exactly for this reason that the valence angle HNH (107.8 deg) is close to a tetrahedral one. The slight deviation of this angle from 109.5 deg is explained, as for the water molecule, by asymmetry in the distribution of the electron clouds about the nucleus of the nitrogen atom: of four electron

Space structure of molecules

There is no direct relationship between the formula of a compound and the shape of its molecules. The shapes of these molecules can be predicted from their Lewis structures, however, with a model developed about 30 years ago, known as thevalence-shell electron-pair repulsion (VSEPR) theory.

The VSEPR theory assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom

Molecular Orbital Method

Consider again the approach of two isolated atoms.  As they get closer together, the orbitals of the outer electron shell of one atom start to overlap with the orbitals of other atoms. An electron in the region of the overlap comes under the influence  of both nuclei and cannot be said to be in an `atomic orbital` any longer: instead it is in a molecular orbital.

The general principle is called linear combination of atomic orbitals (L.C.A.O.):The overlap of  N atomic orbitals produces N molecular orbitals.The average energy of the M.O.`s produced by L.C.A.O. is the same as the average  energy of the overlapping A.O. The M.O.`s are filled with electrons in accordance with the aufbau and Hund principles.

Consider example of  two hydrogen atoms, 1s1 and 1s1.

The two molecular orbitals that different both in character and in energy forms. One is split into two volumes spreading out and away from the nuclei;  the other covers mainly the volume of space between the nuclei. Their average  energy equals of the 1s-atomic orbitals.

Thebonding M.O. is the one that occupies the volume of space between the nuclei. Electron density in this orbital leads to the bonding effect.

Thesplit-volume M.O. is anti-bonding that result in repulsion and the tendency for the system to revert to separate atoms.

It is called asigma (s) molecular orbital because it looks like ans orbital when viewed along the H-H bond. Electrons placed in the other orbital spend most of their time away from the region between the two nuclei. This orbital is therefore anantibonding, orsigma star (s*), molecular orbital.

The M.O.`s are filled  according to the aufbau principle.

Electrons are added to molecular orbitals, one at a time, starting with the lowest energy molecular orbital. The two electrons associated with a pair of hydrogen atoms are placed in the lowest energy, or bonding, molecular orbital, as shown in the figure below. This diagram suggests that the energy of an H2 molecule is lower than that of a pair of isolated atoms. As a result, the H2 molecule is more stable than a pair of isolated atoms.

Molecular Orbitals of the Second Energy Level

The 2s orbitals on one atom combine with the 2s orbitals on another to form as2s bonding and as2s* antibonding molecular orbital.

If we arbitrarily define theZ axis of the coordinate system for the O2 molecule as the axis along which the bond forms, the 2pz orbitals on the adjacent atoms will meet head-on to form as2p bonding and as2p* antibonding molecular orbital, as shown in the figure below.

The 2px orbitals on one atom interact with the 2px orbitals on the other to form molecular orbitals that have a different shape, as shown in the figure below. These molecular orbitals are calledpi (p) orbitals because they look likep orbitals when viewed along the bond. Whereassands* orbitals concentrate the electrons along the axis on which the nuclei of the atoms lie,pandp* orbitals concentrate the electrons either above or below this axis.

The 2px atomic orbitals combine to form apx bonding molecular orbital and apx* antibonding molecular orbital. The same thing happens when the 2py orbitals interact, only in this case we get apyand apy* antibonding molecular orbital.

The interaction of four valence atomic orbitals on one atom (2s, 2px, 2py and 2pz) with a set of four atomic orbitals on another atom leads to the formation of a total of eight molecular orbitals:s2s,s2s*,sp,s2p*,px,py,px*, andpy*.

Such  (consistency) order of adding of electrons to M.O. exists for diatomic molecules of the second period (B,C,N):

s1s <s1s*<s2s <s2s*<px <py <s2p <px*<py*<s2p*.

This order changes for remaining elements of the second period (O, F, Ne):

s1s <s1s*<s2s <s2s*<s2p <px <py <px*<py*<s2p*.

Because they meethead-on, the interaction between the 2pz orbitals is stronger than the interaction between the 2px or 2py orbitals, which meet edge-on. As a result, thes2p orbital lies at a lower energy than thepx andpy orbitals, and thes2p* orbital lies at higher energy than thepx* andpy* orbitals, as shown in the figure below.

Unfortunately an interaction is missing from this model. It is possible for the 2s orbital on one atom to interact with the 2pz orbital on the other. This interaction introduces an element ofs-p mixing, or hybridization, into the molecular orbital theory. The result is a slight change in the relative energies of the molecular orbitals, to give the diagram shown in the figure below. Experiments have shown that O2 and F2 are best described by the model in the figure above, but B2, C2, and N2 are best described by a model that includes hybridization, as shown in the figure below.

The idea ofhybridization andhybrid orbitals comes from the attempt to make overlap theory constituent with the observed shape of molecules. By assuming that  the strength of a bond is directly proportional to the amount of overlap that take place,  it is possible to find more favourable solution which defines an alternative set of orbitals. These orbitals have higher energy level than s-, p-, d- of corresponding shell. The increased overlap causes a larger drop in the energy of the system so that the whole process becomes energetically favourable despite the extra energy needed initially for the hybrid orbitals to be filled. The M.O.`s formed from the overlap of the hydrogen 1s-orbital in the two possible cases are compared below.

Even though the sp3-orbitals is of higher energy than the 2p-orbital, its occupation leads to a lower energy M.O.

Diatomic molecules that consist of  atoms of different sorts (AB) have different values of initial atomic orbitals. Bonding orbitals have energy closer to the energy of the orbitals of the most electronegative atoms, anti-bonding - least electronegative one. Difference in energy initial atomic orbitals define polarity of the bond. Consider the CO molecule as example. Existence of 6 bonding and 0 anti-bonding electrons as in N2 molecule correspond formation of the triple bond. Similarity of such physical properties as melting boiling points, structure and etc.

Bond Order

The number of bonds between a pair of atoms is called thebond order. Oxygen, for example, has a bond order of two.

In molecular orbital theory, we calculate bond orders by noting that there are eight valence electrons in bonding molecular orbitals and four valence electrons in antibonding molecular orbitals in the electron configuration of this molecule. Thus, the bond order is two.

Although the Lewis structure and molecular orbital models of oxygen yield the same bond order, there is an important difference between these models. The electrons in the Lewis structure are all paired, but there are two unpaired electrons in the molecular orbital description of the molecule. As a result, we can test the predictions of these theories by studying the effect of a magnetic field on oxygen.

Atoms or molecules in which the electrons are paired arediamagnetic repelled by both poles of a magnetic. Those that have one or more unpaired electrons areparamagnetic attracted to a magnetic field. Liquid oxygen is attracted to a magnetic field and can actually bridge the gap between the poles of a horseshoe magnet. The molecular orbital model of O2 is therefore superior to the valence-bond model, which cannot explain this property of oxygen.

Using the Molecular Orbital Model to Explain Why Some Molecules Do Not Exist

This molecular orbital model can be used to explain why He2 molecules don't exist. Combining a pair of helium atoms with 1s2 electron configurations would produce a molecule with a pair of electrons in both the bonding and the * antibonding molecular orbitals. The total energy of an He2 molecule would be essentially the same as the energy of a pair of isolated helium atoms. The net effect is non-bonding: there is no net  energy drop as a result of the electrons in the M.O.`s formed.

The molecular orbital diagram for an O2 molecule would therefore ignore the 1s electrons on both oxygen atoms and concentrate on the interactions between the 2s and 2p valence orbitals.

Intermolecular Interaction

The forces retaining the particles of a liquid or solid near one another are of an electrical nature. But these forces differ quite appreciably depending on what the particles are—whether they are atoms of a metal or non-metal element, ions, or molecules.

Intermolecular interaction occurs in substances with a molecular structure. The forces of intermolecular interaction, also known asvan-der-Waals forces, are weaker than covalent forces, but manifest themselves at greater distances.

If a substance consists of polar molecules, for instance, molecules of H2O or HCl, in the condensed state neighbouring molecular dipoles are oriented relative to each other with oppositely charged poles, and as a result their mutual attraction is observed. This kind of intermolecular interaction is called orientational or dipole-dipoleinteraction. The thermal motion of the molecules prevents their mutual orientation.

In substances consisting of non-polar molecules, but that are capable of polarization, for instance CO2, induced dipoles appear. The result is the mutual attraction of the molecules. Thisinduction interaction is also observed in substances with polar molecules, but it is usually considerably weaker than dipole-dipole interaction.

Finally, the motion of the electrons in atoms, and also the vibration of the nuclei and the associated continuous change in the mutual position of electrons and nuclei lead to the appearance ofinstantaneous dipoles. As shown by quantum mechanics, instantaneous dipoles appear in solids and liquids in agreement, the closest portions of neighbouring molecules next to each other are oppositely charged, which leads to their attraction. This phenomenon, calleddispersion interaction, occurs in all substances in the condensed state. Particularly, it underlies the transition of the noble gases at low temperatures into the liquid state.

The weak force of attraction between the neighbouring oppositely charged ends of two instantaneous dipoles is called a van der Waals force.

The greater the number of electrons in a particle, the more likely the electrons are to be momentarily localized at one part of the particle in preference to another. The strength of the van der Waals' force therefore increases with electron number.

Hydrogen Bond

A hydrogen bond is the electrostatic force of attraction between the poorly shielded proton of a hydrogen atom bonded to a small highly electronegative atom such as nitrogen, oxygen, or fluorine, and the lone pair of a neighbouring molecule

The appearance of a hydrogen bond can be explained in a first approximation by the action of electrostatic forces. Thus, in the formation of a polar covalent bond between a hydrogen atom and a fluorine atom, which is highly electronegative, the electron cloud originally belonging to the hydrogen atom is greatly displaced towards the fluorine atom. As a result, the fluorine atom acquires a considerable effective negative charge, while the nucleus of the hydrogen atom (proton) at the "external" side relative to the fluorine atom is almost deprived of an electron cloud. Electrostatic attraction between the proton of the hydrogen atom and the negatively charged fluorine atom leads to the formation of a hydrogen bond. The reason is that a hydrogen ion (proton) is capable of penetrating into the electron shells of other atoms because it has negligibly small dimensions and, unlike other cations, has no inner electron layers that are repelled by negatively charged atoms.

The formation of a hydrogen bond when two HF molecules interact can be represented as follows:

Here the dotted line designates a hydrogen bond, and the symbols "+" and "—" relate to the effective charges of the atoms.

It is well known that compounds in which the hydrogen atom is directly bound to atoms of fluorine, oxygen, and nitrogen, have a number of anomalous properties. For example, the melting points in the series HCl-HBr-HI are —114.2, —86.9, and —50.8°C, respectively. A similar relationship is observed in the series HgS-HzSe-HgTe. As shown byFigs. 11 and 12, however, hydrogen fluoride and water melt and boil at anomalously high temperatures.

The energy of a hydrogen bond is considerably lower than that of an ordinary covalent bond (150 to 400 kJ/molThis energy, however, is sufficient to cause associationof molecules, i. e. their combinationinto dimers(doubled molecules) or polymers, which sometimes exist not only in the liquid state of a substance, but are also retained when it transforms into a vapour. It is the reason for the anomalously high melting and boiling points of substances such ashydrogen fluoride, water, and ammonia.

The hydrogen bond is the reason of certain important features of water—a substance playing a tremendous role in processes occurring in animate and inanimate nature. It also determines to a considerable extent the properties of such biologically important substances asproteins and nucleic acids.


The bands in a metal lattice are either partly filled or, when filled, overlap with other unfilled bands. In both of these cases, there are empty orbitals lying very close to the uppermost filled level. It needs very little energy indeed to excitethe electrons into these empty orbitals. The electrons are therefore extremely mobile and can readily move through the lattice when excited either by electrical or thermal energy.

The electrical conductivity of a metal is the result of the concerted movement of delocalised electrons through the lattice. The thermal conductivity of a metals the result of the exchange of kinetic energy between the delocalized electrons in the lattice.

Unlike metals, nonmetals do not usually have an appreciable electronic conductance. The conduction band is separated from the valence one bya forbidden band, i.e. by a considerable energy gapDE.

In dielectrics, the electrons cannot freely migrate along a crystal and be carriers of an electric current.

Semiconductrs have specific properties distinguishing them from both metals and dielectrics. At low temperatures, their electrical resistance is quite high, and in these conditions they display properties of dielectrics. But when heated or illuminated, the electrical conductance of semiconductors sharply grows and may reach values comparable with the conductance of metals.

The dependence of the electrical properties of semiconductors on the temperature and illumination is explained by the electron structure of their crystals. Here, as in dielectrics, the valence band is separated from the conduction band by the forbidden band Fig. 34, semiconductor).But the width of the forbidden bandDE is not great for semiconductors. Consequently, under the action of quanta of radiant energy or when heated, the electrons occupying the upper levels of the valence band can pass into the conduction band and carry an electric current. With elevation of the temperature or with a growth in the illumination, the number of electrons passing into the conduction band increases; accordingly, the electrical conductance of the semiconductor grows.

When electrons pass over to the conduction band, energy levels not completely occupied by electrons appear in the valence bands. They are known as electron vacancies or "holes". In an electric field, the holes behave like positive electric charges. Hence, a current can be carried in semiconductors both by the electrons of the conduction band (n-conductivity, from the word negative), and by the holes of the valence band (p-conductivity, from the word positive).

When the hydrogen molecule forms from two hydrogen atoms the interatomic distance is less than the sum of two radii of hydrogen atoms (2x0.053 nm). It equals 0.074 nm. It is evident that overlapping of electron clouds of hydrogen atoms characterised the formation of covalent bond.

When the bonds forms two energy processoccur: first one  - is expenditure of energy (lost) on exciting  state of valence electrons atoms (for example carbon atom) and energy that isolates when bonds forms.

As a rule if energy that spend when the bonds forms  is higher than expedinture of energy  on transition of atom to the exciting state.


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