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The origin and nature of chemical bonds

Лекция

Химия и фармакология

The bond dissociation energy, also called simply the 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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2015-07-11

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The origin and nature of chemical bonds

When atoms interact, a chemical 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 as chemical 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 simply
the 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 called ionic bonds.  The force of attraction between ions may be expressed by the inverse square law 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 and electron-pairs shared by the bonding atoms. Lewis called an electron-pair bond between atoms a covalent bond. He proposed that chemical bonds form between electrons residing in the outermost, or valence, shell of each bonding atom. The valence electrons are 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 on difference 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 bond in molecules with 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 a metallic 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 bond is 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.

Valence

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 an ammonium 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 acceptor of 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 by the 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.

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 molecules H2, N2, and Cl2). 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 called effective charges of the atoms in  a molecule.

A polar covalent bond in molecule, for example, in a hydrogen chloride molecule  can be considered as a electric dipoles. Although the net charge of a dipole is zero, an electric field is formed in the space surrounding it as shown in Fig. 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 charge q of an electron and the distance l 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.

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 such dipole-dipole interaction is the mutual attraction of polar molecules and strengthening of the bonds between them. This is why substances 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.

Hybridization of atomic electron orbitals

The method of hybridization of atomic orbitals proceeds from the assumption that in the formation of a molecule, instead of the initial atomic s-, p-, and d-electron clouds, equivalent "blended" or hybrid electron 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.

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 as van-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-dipole interaction. 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. This induction 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 of instantaneous 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, called dispersion interaction,

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.

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 association of molecules, i. e. their combination into 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 as hydrogen 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 as proteins and nucleic acids.

Conductivity

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 excite the 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 by a forbidden band, i.e. by a considerable energy gap DE.

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


 

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