Chemical Bond:-
“The attractive force which holds together the constituents particles (atoms, ions, or molecules) in a chemical species is known as Chemical Bond”.
The tendency of an element to combine with one another is directly related to the valency of the elements. The valency is the ability or tendency of elements to combine with one another. Since the valency depends upon the electronic configurations of the atoms.
A number of attempts were made to explain the formation of chemical bonds between atoms in terms electrons. In 1916, Kossel and Lewis succeeded independently in giving first successful explanation about the cause of combination between atoms based upon the understanding of electronic configuration of noble gases. It has been observed that atoms of noble have little or no tendency to combine with each other or with atoms of other elements. This means that these atoms must be having stable electronic configurations. The atoms of all noble gases (with the exception of helium) have eight electrons in their valence shell. Helium on the other hands, has two electrons in its valence shell. The electronic configuration of the valence shell for all noble gas atoms except helium can be expressed as ns2 np6 (for helium, 1s2) and this represents the stable configuration and corresponds to maximum stability. Due to the stable configuration, the noble gas atoms neither have any tendency to gain nor lose electrons and, therefore, their combining capacity (or valency) is zero.
Thus, the tendency or urge of atoms of various elements to attain stable configuration of eight electrons in their valence shells, is the cause of chemical combination. The principle of attaining maximum of eight electrons in the valence shell of atoms is called octet rule.
Limitation of Octet Rule:-
Although octet rule is useful in a large number of cases, it has many exceptions also. Some important exceptions of octet rule are given below:
1) Hydrogen has one electron in its first energy shell (n =1). It needs only one more electron to fill this shell, because the first shell cannot have more than two electrons. This configuration (1s2) is similar to that of noble gas helium and is stable. In this case, therefore, octet is not needed to achieve a stable configuration.
2) The octet rule cannot explain the formation of certain molecules of beryllium, boron, aluminium, etc. (BeH2, Becl2 BH3, and BF3) in which the central atom has less than eight electrons in the valence shell as shown below:
On the basis of octet rule, therefore, the elements of groups IA, IIA and IIIA should not form stable molecules. This is because they cannot achieve the octets by electron sharing. However, these elements form stable molecules.
3) There are many stable molecules which have more than eight electrons in their valence shells. For example, PF5 has ten SF6 has twelve and IF7 has fourteen electrons around the central atoms P, S and I respectively, as shown below: –
4) It maybe noted that the octet rule is based upon the chemical inertness of noble gases. However, it has been found that some noble gases (especially xenon and krypton) also combine with oxygen and fluorine to form a large number of compounds such as XeF2, KrF2, XeOF2, XeOF4, XeF6, etc.
Lewis Symbols:-
All the electrons in an atom are not involved in the process of combination. The inner shell electrons are well protected and, therefore, they generally do not take part. Lewis pictured the atom in terms of positively charged ‘kernel’ (the nucleus plus the inner electrons) and the outer shell that could accommodate a maximum of eight electrons. Thus, it is mainly the electrons present in the outermost shell that take part in chemical combinations. Therefore, these are also called valence shell electrons. Lewis introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols or electron dot symbols.
Lewis symbols or electron dot symbols:-
According to Lewis notations, the symbol of the element represents the whole of the atom except the valence electrons (i.e., nucleus and the electrons in the inner energy shells). The valence electrons are represented by placing dots (•) or crosses (x) around the symbol. For example, the Lewis symbols for the atoms of second period are given below —
Li , Be , B , C , N , O , F , Ne
Similarly, the Lewis symbols for the elements of third period are:
Na , Mg , Al , Si , P , S , Cl , Ar
Significance of Lewis Symbols:-
(i) The Lewis symbols indicate the number of electrons in the outermost or valence shell.
(ii) Lewis symbols help to predict the common valence of the element. For example, lithium has one electron in valence shell and it can involve this electron in chemical combination process (by losing or sharing) and it is therefore, monovalent; beryllium has 2 electrons for participating in chemical combination and is divalent. Similarly, B and C are trivalent and tetravalent because they have three and four electrons respectively. Thus, for Li, Be, B and C, the number of electrons also indicates the common valence of these elements. However, the common valence for N, O, F and Ne is equal to eight minus the valence electrons. For example, for N it is 3, for O it is 2, for F it is one and for Ne it is 0. Thus, the common valence of an element is either equal to number of dots in the Lewis symbol or it is equal to 8 minus the number of dots.
The chemical bonding may be classified into three main classes-
1) Ionic or Electrovalent Bond
2) Covalent Bond and
3) Co-ordinate Covalent Bond
1) Ionic or Electrovalent Bond:- Ionic or Electrovalent Bond is formed by the transference of one or more electrons from one atom to another. This type of bond usually comes into existence between a metal and a non metal atom. The one atom loses its valence electrons and gets converted into a positive ion (cation) while another atom gains electrons and gets converted into negative ion (anion). The oppositely charged ions are held together by electrostatic force of attraction. Therefore, ionic or electrovalent bond is the electrostatic force of attraction holding the oppositely charged ions. The compound containing ionic bonds are called Ionic or Electrovalent compounds.
Let us illustrate the formation of ionic bond by considering the examples of sodium chloride (Nacl). Sodium atom (1s2 2s2 2p6 3s1) has only one electron in the valence shell and by losing this electron, it can acquire stable electronic configuration of neon (1s2 2s2 2p6).
On the other hand, chlorine atom (1s2 2s2 2p6 3s2 3p5) has seven electrons in its valence shell and needs only one electron to complete its octet. Thus, both the atoms can complete their octets if sodium atom gives one electron to chlorine atom. This tendency is responsible for bonding between sodium and chlorine atoms. Therefore, sodium gives one electron and becomes positively charged Na+ ion, while chlorine takes up the electron and becomes negatively charged, C1– ion. These two ions are then held together by electrostatic forces of attraction which constitute the ionic bond. The ionic bond between Na and Cl may be represented as:
Na ——> Na+ + e–
Cl + e– ——> Cl–
Na+ + Cl– ——-> NaCl or NaCl
Ionic compound
Formation of CaF2 :- The electronic configurations of calcium and fluorine atoms are
Ca (Z = 20): [Ar] 4s2, F (Z = 9): 1s2 2s2 2p5
In the formation of calcium fluoride, calcium loses its both the valence electrons to two fluorine atoms each of which is in need of one electron. This results in the formation of Ca2+ ion and two F– ions. Each of these ions acquires noble gas configuration. One Ca2+ and two F– ions form bonds to give calcium fluoride. This may be represented as:
Ca ———> Ca2+ + 2e–
F + e ——-> F–
Ca2+ + 2F– ——-> CaF2
In the formation of electrovalent bond, the number of electrons lost or gained by an atom is called its electrovalence or electrovalency. It is also equal to the number of unit charges on the ion. For example, sodium is assigned a positive electrovalence of one; calcium is assigned a positive electrovalence of two. Similarly, chlorine and fluorine both are assigned a negative electrovalence of one. Atoms that readily lose electrons are called electropositive while those which readily gain electrons are called electronegative.
Properties of Ionic compound:-
Some of the important characteristics of ionic compound are as follows:
1) The ionic compounds are hard and rigid because of the oppositely charged ions are tightly bound together by strong attractive forces.
2) Ionic compounds have high melting point and boiling points due to the strong binding force.
3) Ionic compounds conducts electricity either in the fused state or in aqueous solution.
4) Most of the ionic compounds are soluble in polar solvent like water.
Note: Formation of ionic bond is affect by the following factors:
- Low ionization energy and high electron affinity.
- Large amount of lattice energy.
Lattice Energy: The stability of ionic compound is determined in terms of their lattice energy. Lattice energy may be defined as – the amount of energy released when one mole of ionic crystal is formed from its constituent ions in the gaseous state. For example,
M+(g) + X–(g) ——— MX (s) – ∆H ,
Where,
∆H = Lattice Energy.
Covalent Bond:-
The bond formed by mutual sharing of electrons between the combining atoms of the same or different elements is called a covalent bond. Therefore, covalent bond is the force of attraction which arises by the mutual sharing of electrons between the atoms. The compounds, thus, formed are called covalent compounds.
Let us consider the formation of a hydrogen molecule from two hydrogen atoms. When two hydrogen atoms share a common pair of electrons between them each of atom attains the stable electronic configuration of the nearest noble gas configuration, helium. This may be represented as:
H• + H• ———- H : H or H—H
Similarly, two chlorine atoms combine with each other to form a molecule of chlorine. In this case, both the atoms have 7 electrons in the outermost shell and have one electron less than argon configuration. Therefore, they contribute one electron each to form a shared pair between two atoms. In this process, both chlorine atoms attain the stable electronic configuration of noble gas (octet). This may be depicted as:
Cl + Cl —— Cl : Cl or Cl—Cl
Multiple covalent bonds:-
When the atoms share one electron pair, the bond formed is called single covalent bond. However, if two electron pairs are shared by the atoms, the bond formed is called double covalent bond.
׃O׃ + ׃O׃ ——- ׃O ׃׃ O O=O
Similarly, when the atoms share three electron pairs the bond is called triple covalent bond.
N + N ——— N N or N≡N
The double and triple covalent bonds are collectively called multiple covalent bonds. Let us now study the examples for double and triple bonds.
Properties of Covalent compound:-
Followings are the characteristic properties of covalent compound-
1) Covalent compounds have low melting and boiling point i.e. they are volatile in nature.
2) Covalent compounds are bad conductor of electricity and carry no current in the fused or aqueous medium.
3) Covalent compounds are insoluble in polar solvent but they dissolves in organic solvent (non-polar) such as benzene, chloroform, acetone, ether, etc.
4) Covalent compound exhibits isomerism i.e. same molecular formula give two or more than two compounds.
Co-ordinate Covalent Bond:-
Co-ordinate Covalent Bond is formed by mutual sharing of electrons between the two atoms but the shared electron-pair is contributed only by one of the two atoms, the other atom simply participates in sharing. The atom which donates an electron pair for sharing is called donor and it must have already completed its octet. On the other hand, the atom which accepts the electron pair in order to complete its octet is called acceptor. The bond is represented by an arrow pointing from the donor towards acceptor. For example, formation of ammonium ion (NH4+), Ozone molecule (O3), Hydronium ion (H3O+), SO2 molecule, etc.
Polar Character of Covalent Bond:-
When covalent bond is formed between two similar atoms, the shared pair of electrons lies midway between the nuclei of the two atoms, because both the atoms have same attraction for the bonding electrons. The molecular orbital constituting the covalent bond is symmetrically distributed around the atoms. Such a covalent bond is called a non-polar covalent bond.
A : A A—A
Symmetrical electron cloud Non-polar bond
For example, molecules like H2, O2, Cl2 and N2 contain non-polar bonds.
In case, the covalent bond is formed between two dissimilar atoms, one of which has a larger value of electronegativity, the bonding pair of electrons is displaced towards the more electronegative atom. In other words, electron cloud containing the bonding electrons gets distorted and the charge density concentrates around the more electronegative atom.
Due to the unequal distribution of electron charge density, the more electronegative atom acquires a partial negative charge indicated as δ– whereas the less electronegative atom acquires a partial positive charge indicated as δ+. Thus, a covalent bond develops a partial ionic character as a result of the difference of electronegativities of the atoms comprising the bond. Such a bond is called polar covalent bond as shown below:
For example, the bond between hydrogen and chlorine atoms in HC1 molecule is polar because the shared electron pair is displaced toward chlorine atom which is more electronegative.
H: Cl or H —->—– Cl or H—Cl
The extent of ionic character in a covalent bond depends upon the difference of electronegativities of the two atoms forming a bond. Greater the difference of electronegativities, greater is the percentage of ionic character in a bond. It has been observed that the bond has 50% ionic character and 50% covalent character if the difference of electronegativities of the participating atoms is 1.7. On the other hand, the covalent character dominates if the difference of electronegativities is less than 1.7 while ionic character dominates if the difference of electronegativities is greater than 1.7. A bond having ionic character more than 50% is usually called ionic bond while that having ionic character less than 50% is called covalent bond.
Dipole Moment:-
It has been observed that a covalent bond between two atoms acquires a partial polar character if the values of electronegativity of the two bonded atoms differ. The two charged ends of the bond behave as electrical dipole and the degree of polarity is measured in terms of dipole moment. Dipole moment is defined as the product of the magnitude of charge on any one of the atoms and the distance between them. Dipole moment is represented by a Greek letter ‘µ’. It can be expressed mathematically, as
µ= e x d
Where, e = charge on any one of the atoms
d = distance between the atoms.
Since the charge e’ is of the order of 10-10 e.s. u. and ‘d’ is of the order of 10-8 cm, then µ, which is the product of ‘d’ and ‘e’ is of the order of 10-18, e.s. u. -cm. This unit is called Debye and is represented by D.
In SI units dipole moment may be expressed in terms of Coulomb-meter abbreviated as Cm.
1 D = 1 x 10-18 e.s.u. = 3.335 x 10-30 C m
For example, dipole moment of HCl is 1.03 x 10-18 e.s.u.-cm and is expressed as 1.03 D.
Covalent Character in Ionic Bond:-
A covalent bond develops ionic character due to the difference of electronegativities of bonded atoms, the ionic bond also develops covalent character when two oppositely charged ions come close, the positive ion tends to distort the electron cloud of negative ion towards itself. Consequently, the electron cloud of negative ion gets polarised and electron density is pulled in between the nuclei of the two atoms. This is precisely what happens in the formation of a covalent bond, i.e., build up of electron density in between the two nuclei as shown. In other words, the ionic bond does not remain 100% ionic but develops some covalent character.
Fig: Distortion of electron cloud of anion.
Note:-
- The power of cation to cause distortion in the electron cloud of -ve ion is referred to as its polarishig power.
- The ability of anion to undergo distortion is called its polarisability.
- The extent of covalent character in ionic bond depends on the polarising power of cation and polarisability of anion which are decided on the set of rules called Fajan’s Rule.
(a) Smaller the size of cation, larger will be its polarising power. For example, Li+ is smaller than K+ ion. Therefore, LiCl has more covalent character than KCl.
(b) For two cations of similar size, the one with pseudo noble gas configuration ns2np6nd10 has larger polarising power than the one with noble gas configuration. For example, CuCl is more covalent than NaC1 because polarising power of Cu+ ion which has pseudo noble gas configuration is more than Na+ ion.
(c) Larger the size of anion, higher will be its polarisability. For example, LiI is more covalent than LiF. Similarly, AIF3 is ionic but AlCl3 is covalent in nature.