Class XI, CHEMISTRY, "Chemical Bond"

Chemical Bond
Introduction
Atoms of all the elements except noble gases have incomplete outermost orbits and tends to complete them by chemical combination with the other atoms.
In 1916, W Kossel described the ionic bond which is formed by the transfer of electron from one atom to another and also in 1916 G.N Lewis described about the formation of covalent bond which is formed by the mutual sharing of electrons between two atoms.
Both these scientists based their ideas on the fact that atoms greatest stability when they acquire an inert gas electronic configuration.
Definition
When two or more than two atoms are combined with each other in order to complete their octet a link between them is produced which is known as chemical bond.

OR
The force of attraction which holds atoms together in the molecule of a compound is called chemical bond.
Types of Chemical Bond
There are three main types of chemical bond.
1. Ionic bond or electrovalent bond
2. Covalent bond
3. Co-ordinate covalent bond or Dative covalent bond
Ionic Bond OR Electrovalent Bond
Definition
A chemical bond which is formed by the complete shifting of electron between two atoms is called ionic bond or electrovalent bond.

OR
The electrostatic attraction between positive and negative ions is called ionic bond.
Conditions for the Ionic Bond Formation
1. Electronegativity
Ionic bond is formed between the element having a difference of electronegativity more than 1.7 or equal to 1.7 eV.
Therefore ionic bond is generally formed between metals (low electronegative) and non-metal (high electronegative) elements.
2. Ionization Potential
We know that ionic bond is formed by the transference of electron from one atom to another, so in the formation of ionic bond an element is required which can lose its electrons from the outer most shell. It is possible to remove electron from the outermost shell of metals because of their low ionization potential values.
3. Electron Affinity
In the formation of ionic bond an element is also required which can gain an element is also required which can gain electron, since non-metals can attract electrons with a greater force due to high electronegativity. So a non-metal is also involved in the formation of ionic bond due to high electron affinity.
Example of Ionic Bond
In order to understand ionic bond consider the example of NaCl. During the formation of Ionic bond between Na and Cl2, Sodium loses one electron to form Na+ ion while chlorine atom gains this electron to form Cl- ion. When Na+ ion and Cl- ion attract to each other NaCl is formed. The stability of NaCl is due to the decrease in the energy. These energy change which are involved in the formation of ionic bond between Na and Cl are as follows.
i. Sodium has one valence electron. In order to complete its octet Na loses its valence electron. The loss of the valence electron required 495 kJ/mole.
Na —-> Na+ + e- ………………….. ΔH = 495 kJ/mole
ii. Chlorine atom has seven electrons in its valence shell. It require only one electron to complete its octet, so chlorine gains this electron of sodium and release 348 kJ/mole energy.
Cl + e- —-> Cl- …………………. ΔH = -348 kJ/mole
Here the energy difference is 147 kJ/mole (495 – 348 = 147). This loss of energy is balanced when oppositely charged ions are associated to form a crystal lattice.
iii. In third step, positively charged Na+ ion and negatively charged Cl- ion attract to each other and a crystal lattice is formed with a definite pattern.
Na+(g) + Cl-(g) —-> Na+Cl- ……….. ΔH = – 788 kJ/mole
This energy which is released when one mole of gaseous ions arrange themselves in definite pattern to form lattice is called lattice energy.
From this example, we can conclude that it is essential for the formation of ionic bond that the sum of energies released in the second and third steps must be greater than the energy required for the first step.
Characteristics of Ionic Compounds
1. An ionic compounds, the oppositely charged ions are tightly packed with each other, so these compounds exist in solid state.
2. Due to strong attractive forces between ions a larger amount of energy is required to melt or to boil the compound and hence the melting and boiling point of the ionic compound are generally high.
3. Ionic compounds are soluble in water but insoluble in organic solvents like benzene, CCl4. etc.
4. In the aqueous solution, the ionic compounds are good electrolytes, because in water the interionic forces are so weakened that the ions are separated and free to move under the influence of electric current. Due to this free movement of ions, the ionic compounds conduct electricity in their solutions.
Covalent Bond
Definition
A link which is formed by the mutual sharing of electrons between two atoms is called covalent bond.
Explanation
In the formation of covalent bond, mutual sharing of electron takes place. This mutual sharing is possible in non-metals, therefore covalent bond is generally formed between the atoms of non-metals. For example
In Cl2 molecule, two atoms of chlorine are combined with each other to form Cl2 molecule. Each atom of chlorine having seven electrons in its valencies shell. These atoms are united with each other by sharing one of its valence electron as shown.
Cl Cl: —-> :Cl :Cl OR Cl – Cl
In this molecule, one shared pair of electrons forms a single covalent bond between two chlorine the atoms. With the formation of a covalent bond the energy of the system is also decreased.
Cl + Cl —-> Cl – Cl ………….. ΔH = – 242 kJ / mole
This released energy lowered the energy of the molecule and the stability of the compound is also increased.
Types of Covalent Bond
There are three main types of covalent bond.
1. Single Covalent Bond
When a covalent bond is formed by sharing of one electron from each atom, that it is called single covalent bond and denoted by (-) single line between the two bonded atoms e.g.
Cl – Cl, H – H, H – Br etc.
2. Double Covalent Bond
In a covalent bond, if two electrons are shared from each of the bonded atom then this covalent bond is called double covalent bond and denoted by (=) two lines e.g.
O = O, O : : O
3. Triple Covalent Bond
When a covalent bond is formed by sharing of three electrons from each atom then this type of covalent bond is called triple covalent bond, and denoted by (≡) three lines between the two bonded atoms e.g.
N : : N :, N ≡ N
The bond distance of multiple bonds are shorter and the bond energies are higher.
Characteristics of Covalent Compounds
The main characteristics properties of covalent compounds are as follows
1. The covalent compounds exist as separate covalent molecules, because the particles are electrically neutral so they passes solid, liquid or gaseous state. This intermolecular force of attraction among the molecules.
2. Since the covalent compound exist in all the three states of matter so their melting points and boiling point may be high or low.
3. Covalent compounds are non-electrolytes so they do not conduct electricity from their aqueous solution.
4. Covalent compounds are generally insoluble in water and similar polar solvent but soluble in the organic solvents.
Co-Ordinate OR Dative Covalent Bond
Definition
It is a type of covalent bond in which both the shared electrons are donated only be one atom, this type is called co-ordinate covalent bond.

The ∞ ordinate covalent bond between two atoms is denoted by an arrow (→). The atom which donates an electron pair is called as a donor of electron and the other atom involved in this bond is called acceptor. E.g.
A + B —-> A : B OR A → B
Dipole Moment
Definition
The product of the charge and the distance present in a polar molecules is called dipole moment and represented by μ.
OR
The extent of tendency of a molecule to be oriented under the influence of an electric field is called dipole moment.
Mathematical Representation of Dipole Moment
Suppose the charge present on a polar molecule is denoted by e and the separation between the two oppositely charged poles of the molecules is d, then the product of these two may be written as
e x d = μ
Where μ is dipole moment.
Dipole Moment in Diatomic Molecules
The diatomic molecules which are made up of similar atoms will be non-polar and their dipole moment is zero but the diatomic molecules made up of two different atoms e.g. HCl or Hl are polar and have some dipole moment. The value of the dipole moment depends upon the difference of electronegativities of the two bonded atom. If the difference of electronegativity between the atoms is greater, the polarity and also the dipole moment of the molecule is greater e.g.
The dipole moment of HCl = 1.03 debye
Whereas dipole moment of HF = 1.90 debye
Dipole Moment of Poly Atomic Molecules
In poly atomic molecules, the dipole moment of molecules depends upon the polarity of the bond as well as the geometry of the molecule.
Ionic Character of Covalent Bond
In homonuclear diatomic molecules like Cl2, O2, l2, H2 both the atoms are identical so the shared electrons are equally attracted due to identical electronegativities and hence the molecules are non-polar.
When two dissimilar atoms are linked by a covalent bond the shared electrons are not attracted equally by the two bonded atoms. Due to unsymmetrical distribution of electrons one end of the molecules acquire partial positive charge and the other end acquire a partial negative charge. This character of a covalent bond is called Ionic character of a covalent bond.
The ionic character of a covalent bond depends upon the difference of electronegativity of the two dissimilar atoms joined with each other in a covalent bond. E.g., the H-F bond is 43% ionic whereas the H-Cl bond is 17% ionic. The ionic character greatly affects the properties of a molecules e.g., melting point, boiling point of polar molecules are high and they are soluble in polar solvent like H2O. Similarly the presence of partial polar character shortens the covalent bond and increases the bond energies.
Bond Energy
Definition
The amount of energy required to break a bond between two atoms in a diatomic molecule is known as Bond Energy.
OR
The energy released in forming a bond from the free atoms is also known as Bond Energy.
It is expressed in kilo Joules per mole or kCal/mole.
Examples
i. The bond energy for hydrogen molecule is
H – H(g) —-> 2 H(g) …………………….. ΔH = 435 kJ/mole
OR
H(g) + H(g) —-> H – H ………………….. ΔH = 435 kJ/mole
It can be observed from this example that the breaking of bond is endothermic whereas the formation of the bond is exothermic.
ii. The bond energy for oxygen molecule is
O = O(g) —-> 2 O(g) …………………… ΔH = 498 kJ/mole
OR
O(g) + O(g) —-> O = O ……………….. ΔH = -498 kJ/mole
Bond energy of a molecule also measure the strength of the bond. Generally bond energies of polar bond are greater than pure covalent bond.
E.g.
Cl – Cl —-> 2 Cl …………………… ΔH = 244 kJ/mole
H – Cl —-> H+ + Cl- ………………. ΔH = 431 kJ/mole
The value of bond energy e.g., triple bonds are usually shorter than the double bond therefore the bond energy for triple bond is greater than double bond.
Sigma & PI Bond
Sigma Bond Definition
When the two orbitals which are involved in a covalent bond are symmetric about an axis, then the bond formed between these orbitals is called Sigma Bond.

OR
A bond which is formed by head to head overlap of atomic orbitals is called Sigma Bond.
Explanation
In the formation of a sigma bond the atomic orbital lies on the same axis and the overlapping of these orbital is maximum therefore, all such bonds, in which regions of highest density around the bond axis are termed as sigma bond.
Types of Overlapping in Sigma Bond
There are three types of overlapping in the formation of sigma bond.
1. s-s orbitals overlapping
2. s-p orbitals overlapping
3. p-p orbitals overlapping
In all the three types, when the two atomic orbitals are overlapped with each other two molecular orbitals are formed. In these two molecular orbitals the energy of one orbital is greater than the the atomic orbitals which is known as sigma antibonding orbital while the energy of the other orbital is less than the atomic orbital this orbital of lower energy is called sigma bonding orbital and the shared electron are always present in the sigma bonding orbitals.
1. s-s Orbitals Overlapping
In order to explain s-s overlapping consider the example of H2 molecule. In this molecule is orbital of one hydrogen overlaps with is orbital of other hydrogen to form sigma bonding orbitals. Due to this bonding a single covalent bond is formed between the two hydrogen atoms.
Diagram Coming Soon
2. s-p Orbitals Overlapping
This type of overlapping takes place in H-Cl molecule. 1s orbital of hydrogen overlaps with 1p orbital of chlorine to form a single covalent bond. In this overlapping two molecular orbitals are formed, one of the lower energy while the other orbital is of higher energy. The shapes of these orbitals are as follows.
Diagram Coming Soon
3. p-p Orbitals Overlapping
This type of overlapping takes place in fluorine molecule. In this mole 1p orbital of a fluorine atom is overlapped with 1p orbital of the other fluorine atom. The molecular orbitals formed in this overlapping are given in figure
Diagram Coming Soon
PI Bond
When the two atomic orbital involved in a covalent bond are parallel to each other then the bond formed between them is called pi bond.

In this overlapping, two molecular orbitals are also formed. The lower energy molecular orbitals is called π bonding orbital while the higher energy molecular orbital is called π antibonding orbital. The shape of these molecular orbitals are as follows.
Diagram Coming Soon
Hybridization
Definition
The process in which atomic orbitals of different energy and shape are mixed together to form new set of equivalent orbitals of the same energy and same shape.

There are many different types of orbital hybridization but we will discuss here only three main types.
1. sp3 Hybridization
The mixing of one s and three p orbitals to form four equivalent sp3 hybrid orbitals is called sp3 hybridization. These sp3 orbitals are directed from the center of a regular tetrahedron to its four corners. The angles between tetrahedrally arranged orbitals are 109.5º.
It has two partially filled 2p orbitals which indicate that it is divalent, but carbon behaves as tetravalent in most of its compounds. It is only possible if one electron from 2s orbital is promoted to an empty 2pz orbital to get four equivalent sp3 hybridized orbitals.
Diagram Coming Soon
The four sp3 hybrid orbitals of the carbon atom overlap with 1s orbitals of four hydrogen atoms to form a methane CH4 molecule.
The methane molecule contains four sigma bonds and each H-C-H bond angle is 109.5º.
2. sp2 Hybridization
The mixing of one s and two p orbitals to form three orbitals of equal energy is called sp2 or 3sp2 hybridization. Each sp2 orbital consists of s and p in the ratio of 1:2. These three orbitals are co-planar and at 120º angle as shown
Diagram Coming Soon
A typical example of this type of hybridization is of ethane molecule. In ethylene, two sp2 hybrid orbitals of each carbon atom share and overlap with 1s orbitals of two hydrogen atoms to form two σ bonds. While the remaining sp2 orbital on each carbon atom overlaps to form a σ bond. The remaining two unhybridized p orbitals (one of each) are parallel and perpendicular to the axis joining the two carbon nuclei. These generates a parallel overlap and results in the formation of 2 π orbitals. Thus a molecule of ethylene contain five σ bonds and one π bond.
Diagram Coming Soon
3. sp Hybridization
When one s and one p orbitals combine to give two hybrid orbitals the process is called sp hybridization. The sp hybrid orbitals has two lobes, one with greater extension in shape than the other and the lobes are at an angle of 180º from each other. It means that the axis of the two orbitals form a single straight line as shown.
Now consider the formation of acetylene molecule HC ≡ CH. The two C-H σ bonds are formed due to sp-s overlap and a triple bond between two carbon atoms consist of a σ bond and two π bond. The sigma bond is due to sp-sp overlap whereas π bonds are formed as a result of parallel overlap between the unhybridized four 2p orbitals of the two carbon.
Diagram Coming Soon
Valence Shell Electron Pair Repulsion Theory
The covalent bonds are directed in space to give definite shapes to the molecules. The electrons pairs forming the bonds are distributed in space around the central atom along definite directions. The shared electron pairs as well as the lone pair of electrons are responsible for the shape of molecules.
Sidwick and Powell in 1940 pointed out that the shapes of the molecules could be explained on the basis of electron pairs present in the outermost shell of the central atom. Pairs of electrons around the central atom are arranged in space in such a way so that the distances between them are maximum and coulombie repulsion of electronic cloud are minimized.
The known geometries of many molecules based upon measurement of bond angles shows that lone pairs of electrons occupy more space than bonding pairs. The repulsion between electronic pairs in valence shell, decreases in the following order.
Lone Pair – Lone Pair > Lone Pair – Bond Pair > Bond Pair – Bond Pair
When we apply this theory we can see the variation of angle in the molecular structures.
Consider the molecular structures of NH3, OH & H2O.
Diagram Coming Soon
Variation from ideal bond angles are caused by multiple covalent bonds and lone electron pairs both of which require more space than single covalent bonds and therefore cause compression of surrounding bond angles.
Thus the number of pairs of electrons in the valency shell determine the overall molecular shape.
Structure of BeCl2
The two bond pairs of electrons in BeCl2 arrange themselves as far apart as possible in order to minimize the repulsion between them.
Structure of BF3 OR BCl3
In this molecule three bond pair are present around boron to arrange themselves as far apart as possible a trigonal structure is formed.
Hydrogen Bond
When hydrogen is bonded with a highly electronegative element such as nitrogen oxygen, fluorine, the molecule will be polarized and a dipole is produced. The slightly positive hydrogen atom is attracted by the slightly negatively charged electronegative atom. An electrostatic attraction between the neighbouring molecules is set up when the positive pole of one molecule attracts the negative pole of the neighbouring molecule. This type of attractive force which involves hydrogen is known as hydrogen bonding