WBBSE Class 10 Physical Science Notes Chapter 8.2 Ionic and Covalent Bonding

Comprehensive WBBSE Class 10 Physical Science Notes Chapter 8.2 Ionic and Covalent Bonding can help students make connections between concepts.

Ionic and Covalent Bonding Class 10 WBBSE Notes

Chemical Bond: The chemical force that hold the atoms together in a molecule is called the Chemical Bond.

Cause of chemical bonding

1. Tendency to acquire noble gas configuration: From the electronic configuration of the noble gases it is clear that all noble gases (except helium) have eight electrons in their valence shell i.e. they have ns², np⁶configuration.

The noble gas configurations are highly stable and have no tendency to lose or gain electrons. Thus, inert nature of noble gases is due to their stable electronic configuration.

Atoms of all other elements do not have eight electrons in their valence shall (outermost orbit).
Therefore, atoms of these elements combine with other or with atoms of other elements to acquire stable configuration of the nearest noble gas (according to octet rule).

This tendency of atoms of various elements to acquire stable configuration of the nearest noble gas is the cause of chemical combination.

2. Tendency to acquire minimum energy: All systems tend to attain stability by lowering their potential energy. Increase in attractive forces leads to decrease in energy. Conversely, decrease in attractive force or increase in repulsive forces increases energy.

Let us now apply these principles to bonding between atoms. Even, atom bies to attain a state of minimum energy. When two atoms are brought closer to form a molecule, following forces come into play.

Attractive forces between electrons of one atom and nudeus of other. Repulsive forces due to inter-electronic and inter-molecular repulsive forces. The molecule becomes stable when there is a net increase in the attractive force. This Increase in attractive force results in decrease energy. Greater the decrease in energy, more stable the molecule will be.

• Electronic theory of valency: According to electronic theory (developed by Kosel and Lewis. 1916) the valency of on element is the number of electrons that its atoms can gain, lose or share to acquire stable nearest noble gas configuration.

Depending upon the mode of acquiring nearest noble gas configuration, there are three common types of bonds.

• Ionic or Electrovalnet bond
• Covalent bond
• Co-ordinate or Dative bond.

Apart from these chemical bonds, there are some physical bonds (which are electrosintic in nautre). The main types of physical bonds are:

• Hydrogen bond
• Metallic bond
• Vander Waai’s interactions

(a) ionic bonding :

The electrostatic force of attraction which holds the oppositely charged ions together is known as ionic bond or electrovalent bond and the compounds which are formed by the transference of electrons from one atom to the other are called ionic or electrovalent compounds. The number of electrons which an atom loses or gain while forming an ionic bond is known as its electrovalency.

Conditions necessary for the formation of ionic bond

• Formation of cation from a neutral atom having low ionisation energy.
• Formation of an anion from an anion from a neutral atom with high value of electron affinity.
• Formation of crystal lattice from oppositely charged ions involving large release of energy. Higher the Lattic energy of a crystal, more readily it will get formed.

Characteristics of Ionic Compounds:

• All ionic compounds are usually crystalline solids and are composed of ions even in the solid state.
• Ionic solid have high melting points and boiling points due to the presence of strong attractive forces between the oppositely charged ions.
• These are highly soluble in polar solvents (such as water) haveing high dielectric constant but insoluble in organic solvent (such as benzene, alcohol, ether etc.)
• Ionic compounds have low volatility, high density and high stability.
• In molten state or in solution in polar solvents, ionic compounds are good conductors of electricity.
• In solution, ionic compounds undergo ionic reactions which are very fost.
• Crystals of certain Ionic compounds have similar arrangement of atomns as well as geometry.
• Ionic bonds are non-directional and due to the non-directional nature of ionic bonds, ionic compounds do not show isomerism.
• Ionic compounds are hard and brittle

(b) Covalent Bond (Lewis Concept) :

A covalent bond is formed by the mutual sharing of electrons between the atoms, both of which are short of electrons. The compound so formed is called covalent compound. The number of electrons contributed by an atom for sharing is called its covalency. Depending upon the number of electrons shared between two atoms being one, two or three, we have single covalent bond ( : or -), double convalnt bond ( : : or = ) and triple convalent bond ( or ).

Characteristics of Covalent Compounds:

• Covalent compounds exist in solid, liquid or gaseous state.
• Covalent compounds like napthalene in which molecules are held together by weak Vander Waal’s forces, have low melting points and boling points.
• Covakrnt bonds are rigid and directional in nature.
• Covalent solids like diamond, Si, C etc. consisting of giant molecules are bad conductors of electricity due to the absence of free electrons or ions.
• Since i.e dissolves like, polar covalent compounds like sugar are soluble in polar solvents like water and the non-polar compounds like naphthalene are more soluble in non-polar solvents Like benzene.
• Covalent compounds are neither hard nor brittle like ionic compounds.
• Covalent compounds undergo molecular reactions in solution.

Examples of covalent compounds and their structures:

Factors affecting the formation of ionic bonds :

• The ionization energy of the electropositive element should be low i.e. the metal atoms form a low charged positive ion easily.
• The quantitative value of electron affinity of electronegative element should be high i.e. the non-metal atom is small and give rise to low charged negative ion easily.
• Lattice enthalpy should be high.

Lattice enthalpy : The energy given off when gaseous positive and negative ions together to form 1 mole of the solid ionic compound is called Lattice enthalpy (U)
Na⁺ (g) + Cl(g) → NaCl (s)    ΔH = -U
For the reverse process, NaCl (s) → Na⁺(g) + Ch(g)    ΔH = +U
Lattice enthalpy a charge of ions
Lattice enthalpy \(\alpha \frac{1}{\text { size of ions }}\)

Lattice enthalpy play an important role in deciding the solubility of ionic solids.

• If the anion and the cation are of comparable size, the cation radius will influence the lattice energy. Since lattice energy decreases much more than the hydration energy with increasing ionic size, solubility will increase as we go down the group.
• If the anion is large compared to cation, the lattice energy will remain almost constant within a particular group. Since, the hydration energies decrease down a group, solublity will decrease as found for alkaline earth metal carbonates and sulphates.
• Born-Haber cycle : The fundamental of this cycle is based upon the fact

Example :
According to Hesse’s law

Thus, lattice energy of Nacl(S) has a large negative value. This indicates that the compound is highly stable.
Finding the number of covalent bonds.

Finding the number of covalent bond electron shored
Number of covalnent bonds between two atoms = \(\frac{\text { electron shored }}{2}\)
Total number of shared electrons S = N – A

S = the total number of electrons shared in the molecule or polyatomic ion.
N the number of valence shell electrons needed by all the atoms in the molecule or ion to achive noble gas configuration [ N = 8x] number of atoms ( H excluded) + 2 x number of H-atoms]
A = the number of electrons available in the valnece shell of all of the representative atoms. This is equal to the sum of their periodic group number.
e.g. in CO₂ for 0-atom, A = 6 and for C-atom, A = 4
Thus A = 4+2 x 6 = 16
N = 8 x 3=24
S = 24 – 16 = 8
Thus, there are eight electrons shared to form bonds.
: Ö : : C : : Ö :
Formal charge: The atoms of a molecule or ion are usually neutral i.e. carry no charge for some purposes, such as to find reaction mechanism,assigning of formal charge of atoms in a molecule or ion is important.

• Favourable factors for covalent bonding (Fajan’s rule)

(i) The charge on cation or anion must be larger. The increased charge will increases the polarisation of the other ion (anion), thus covalent character is increased.

(ii) The cation must be smaller because in the fact, the charge possess will be more concentrated, thus causing more polarisation of anion. Hence we can say that ionic compounds having smaller cations show more covalent nature.

(iii) The anion must be larger. In such anions, the aoter electrons will be at a greater distance from the nucleus, hence more easily influenced by the attractive forces of cation. As the result larger anions will be more easily polarised in comparison to smaller anions.

(iv) The polarising power of those cations which donot have inert gas configuration will be more in comparision to cations having inert gas configuration.

• Some important covalent bond parameters :

(i) Bond length: The average distance between the centre of nuclei of the two bonded atoms in a molecule is known as bond length. It depends upon the size of atoms hybridization, steric effect, resonance etc.

(ii) Bond enthalpy: Bond enthalpy is the amount of energy required to break a particular bond in one mole of gaseous molecule.
Bond enthalpy α electronegative

(iii) Bond order : It is the number of covalent bonds present between the two atoms in a molecule.
order a bond enthalpya bond length
Bond order a bond enthalpy α \(\frac{1}{\text { bond length }}\)

(iv) Bond angles : It is the angle between the bonded orbitais. Generally it decreases, as the number of lone pair of electrons increases, or as the electronegativity of the central atom decreases.

However, in molecule having same central atom, bond angle increases as the electronegativity of surrounding atom decreases.

∴ The order of eledronegativity is CI > Br > I

• Type of covalent Bonds:

(i) Non-Polar Conalnt Bond : If the covalent bond is formed between two homonuclear atoms i.e. between atoms of exactly equal electronegativity, the electron pair is equally shared between them. Such a bond is called non-polar covalent bond, e.g. H₂, Cl₂, F₂  etc.

(ii) Polar covalent bond : If the bond forming entities are dissimilar i.e. heteronuclear or with different electronegativity, the bond formed has partial ionic character as the electron pair is attracted by more electronegative entity. Such a bond is called polar covalent bond.

The greater difference in electronegativity higher is the polar nature.

Electronegativity : The electronegativity of an element is the tendency or ability its atom to attract the bonding or shared pair of electrons towards itself in a covalent bond. The relative order of electro negativity of some important element is:

Electronegativity is maximum in : F (4.0)
Electronegativity is minimum in: Fr (0.7)

Co-ordinate bond:
A Co-ordinate bond is a special type of covalent bond in which the shared pair of electrons is contributed by one of the two combining atoms. A Co-ordinate bond is represented by an arrow ( – ) pointing from the donor towards the acceptor.
Example:

Characteristics of co-ordinate compounds :

• The compounds exhibit all three states i.e. solid, liquids and gases under ordinary conditions.
• The melting and boiling points of these compounds and higher as compared to covalent compounds but lower than those of ionic compounds.
• Like covalnent compounds these are also poor conduction of electricity in solid as well as fused state.
• These compounds are sparingly soluble in polar solvent like water, however, these are readily soluble in organic solvent.
• Compounds are generally as stable as covalent compounds. The addition compounds are not very stable.
• Co-ordinate linkage is rigid and directional, thus compounds exhibit isomerism.
• Cd-ordinate compounds show molecular reactions just like covalent compounds.
• The dielectric constants of these compounds are higher.

Dipole moment : The product of the magnitude of positive or negative charge (q) and the distance (d) between the centres of positive and negative charges in a polar molecule is called dipole moment.
μ = q x d (p = Dipole moment)

The unit of dipole moment in CGS system in debey and in SI system is Coulomb-metre.

Dipole moment is a vector quantity, thus indicated by the symbol (+→) pointing towards the negative end e.g.

For symmetrical molecules, dipole moment, dipole moment is zero but unsymmetrical molecules have some dipole moment e.g.

However symmetric molecules like HF, H₂O, NH₃, have some dipole moment

By using dipole moment, the precentage ionic character of a polar bond (A-B) is calculated as,

Resonance: The various Lewis structures, which differ in the positions of non-bonding or π-electrons but not in the relative positions of atoms are called resonance structures contributing structure or canonical forms and this concept is known as resonance.

Example : Benzene molecule can be represented as a resonance hybrid (III) of two kekule structures, (I) and (II).

\(\mathrm{BO}=\frac{\text { total number of bonds between two atoms in all the structures }}{\text { total number of canonical forms }}\)

The stability of a resonance structure can be decided by considering the following points-

• A non-polar canonical form is more stable.
• More the number of covalent bonds, more is the stability.
• Resonance structure in which negative charge resides on electronegative atom and positive charge on electropositive atom and positive charge on electropositive atom is more stable as compared to that for which opposite is true.
• Hydrogen bond The electrostatic force of attraction existing between the H-atom covalently bonded to an electronegative atom (F, O or N) in a molecule and the electronegative atom of another molecule (similar or different type) is known as hydrogen bond.

Types of hydrogen bonding

(i) Intermolecular hydrogen-onding : When hydrogen bonding occurs between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding.
Example:

(ii) Intramolecular hydrogen bonding : When hydrogen bonding takes place within the same molecule, it is called intramolecular hydrogen bonding.
Examples :

• Hybridisation : It may be defined as the process of inter mixing of atomic orbitals of the same atom having same or slightly different energies so as to redistribute their energies and give new orbitals of equal energies and identical shapes.

• Types of hybridisation :

Determination of hybridisation of the central atom
H = \(\frac{1}{2} [V+X – C . A ]\)
H = Number of orbitals involved in hybridisation ;
V = Number of electrons in the valence shell of the central atom ;
X = Number of monovalent atoms surrounded the central atom ;
C = Change on cation.
A = Change of anion.

Difference between Electrovalent and Covalent Compounds :

Electrovalent Compound	Covalent Compound
1. These are formed by the transfer of one or more electrons from one atom to another. 2. These consist of ions. 3. These are soluble in water but insoluble in organic solvents. 4. These are hard solids with high melting and boiling points. 5. These conduct electricity in fused as well as aqueous solutions. 6. These undergo ionic reactions which are very fast. 7. These do not show isomerism.	1. These are formed by the sharing of one or more electrons between the bonded atoms. 2. These consist of invidual molecules. 3. These are isoluble in water but solute in organic solvents. 4. These exist as gases, liquids or soft solids with low melting and boiling points. 5. These do not conduct electricity. 6. These undergo molecular reactions which are very slow. 7. These show isomerism.