Maharashtra Board Class 11 Chemistry Chapter 5 Chemical Bonding PDF Download

Chemical Bonding

5.1 Introduction

Why are atoms held together in chemical compounds? There must be some force that holds them together. You have already learned in lower classes that the forces holding atoms together in a compound are the chemical bonds.

How are chemical bonds formed between two atoms? There are two ways of formation of chemical bonds. One is by loss and gain of electrons. The other is by sharing a pair of electrons between the two atoms. In either process of formation of chemical bond each atom attains a stable noble gas electronic configuration.

Which electrons are involved in the formation of chemical bonds? The electrons present in the outermost shell of an atom are involved in the formation of a chemical bond.

5.2 Kossel And Lewis Approach To Chemical Bonding

Number of attempts were made to explain the formation of chemical bond in terms of electrons. But the first satisfactory explanation was given by W. Kossel and G. N. Lewis independently. They gave a logical explanation of valence which was based on the inertness of noble gases. On the basis of this they proposed a theory of valence known as Electronic Theory Of Valence in 1916.

According to Lewis, the atom can be pictured in terms of a positively charged kernel (the nucleus plus inner electrons) and outer shell that can accommodate a maximum of eight electrons. This octet of electrons represents a stable electronic arrangement.

Lewis stated that each atom achieves stable octet during the formation of a chemical bond. In case of sodium and chlorine this can be achieved by transfer of one electron from sodium to chlorine. Thus Na (2, 8) and Cl (2, 8, 8) ions are formed which held together. In case of other molecules like H2, F2, Cl2, HCl etc. the bond is formed by the sharing of a pair of electrons between the atoms. In this process each atom attains a stable outer octet of electrons.

Octet Rule

In 1916 Kossel and Lewis proposed an important theory for explaining the formation of chemical bond known as Electronic Theory Of Valence. This theory is mainly based on octet rule developed by Lewis. Octet rule is based on stability of noble gases due to presence of eight electrons (ns2np6) in the valence shell.

This rule states that during the formation of chemical bond, atom loses, gains or shares electrons so that its outermost orbit (valence shell) contains eight electrons. Therefore the atom attains the nearest inert gas electronic configuration.

The octet rule is found to be very useful in explaining the normal valence of elements and in the study of the chemical combination of atoms leading to the formation of molecule. However it should be noted that octet rule is not valid for H and Li atoms. These atoms tend to have only two electrons in their valence shell similar to that of Helium (1s2) which called duplet.

5.2.1 Ionic Bond

I. Formation Of Sodium Chloride (NaCl)

The electronic configurations of Sodium and Chlorine are:

Na (Z = 11) 1s2 2s2 2p6 3s1 or 2, 8, 1

Cl (Z = 17) 1s2 2s2 2p6 3s2 3p5 or 2, 8, 7

Sodium has one electron in its valence shell. It has a tendency to lose one electron to acquire the configuration of the nearest noble gas Ne (2, 8). Chlorine has seven electrons in its valence shell. It has a tendency to gain one electron and thereby acquire the configuration of the nearest noble gas Ar (2, 8, 8). During the combination of sodium and chlorine atoms, the sodium atom transfers its valence electron to the chlorine atom. So sodium atom changes into Na+ ion while the chlorine atom changes into Cl- ion. The two ions are held together by strong electrostatic force of attraction. The formation of ionic bond between Na and Cl can be shown as follows:

Na + Cl → Na+ + Cl-

2,8,1 2,8,7 2,8 2,8,8

Na+ + Cl- → NaCl or Na+Cl- Ionic bond

Teacher's Note

Atoms lose or gain electrons to become stable and form ionic bonds. For example, when you eat salt in food, the sodium and chlorine atoms are bonded together as NaCl ionic compound.

Exam Trick

Remember: Ionic bond = electrons are transferred from one atom to another. One atom becomes positive (loses electrons) and one becomes negative (gains electrons). Like giving and taking money between two people.

Points to Remember

Ionic bonds form when atoms lose or gain electrons.
Noble gases are very stable because they have a full outer shell.
Sodium loses one electron to become Na+ ion.
Chlorine gains one electron to become Cl- ion.
The attraction between opposite charges holds the ionic bond together.

II. Formation Of Calcium Chloride (CaCl2)

The following representation shows the formation of compound calcium chloride from the elements calcium and chlorine:

Electronic configuration of

Calcium: 1s2 2s2 2p6 3s2 3p6 4s2

Cl: 1s2 2s2 2p6 3s2 3p5

Cl + Ca + Cl → Cl + Ca2+ + Cl

2,8,7 2,8,8,2 2,8,7 2,8,8 2,8,8 2,8,8

Cl + Ca2+ + Cl → CaCl2 or Ca2+(Cl-)

5.1.2 Ionic Solids And Lattice Enthalpy

Ionic solids are solids which contain cations and anions held together by ionic bonds. Kossel treatment helps us to understand the formation of ionic bonds between ions of different elements. Formation of ions depends on the ease with which an atom can lose or gain electrons.

M(g) → M+(g) + e-

X(g) + e- → X-(g) electron gain enthalpy

M+(g) + X-(g) → MX(g) → MX(s)

CsF is the most ionic compound. Because Cs is the most electropositive while F is the most electronegative element. The electronegativity difference between them is the largest. Hence ions are easily separable, the bond is weakest and the compound is least stable ionic compound.

Elements having low ionization enthalpy can readily form ionic bond with elements having a high negative value of electron gain enthalpy. Both these processes take place in gaseous phase. All ionic compounds in the solid state have each cation surrounded by a specific number of anions and vice versa.

Ionization is always an endothermic process while electron gain process can be exothermic or endothermic. Based on the ionization enthalpy (ΔiH) and electron gain enthalpy, we can predict which elements can form ionic compounds.

The arrangements of cations and anions in a crystalline solid is ordered and they are held together by coulombic forces of attraction. During their formation, these compounds crystallize from the gaseous state (MX(g) → MX(s)) to the solid state. The structure in which they crystallize depends upon the size of the ions, their packing arrangement and other factors. The overall stability of the ionic solid depends upon the interactions between all these ions and the energy released during the formation of the crystal lattice.

Let us consider the formation of NaCl ionic solid.

Na(g) → Na+(g) + e- ΔiH = 495.8 kJ mol-1

Cl(g) + e- → Cl-(g) ΔegH - 348.7 kJ mol-1

Na+(g) + Cl-(s) → NaCl(g) + 147.7 kJ mol-1

Conversion of NaCl(g) → NaCl(s) is associated with release of energy which is -788 kJ mol-1. This released energy is much more than the absorbed energy. Thus stability of an ionic compound can be estimated by knowing the amount of energy released during lattice formation and not just by energy associated with completion of octet around the ionic species in the gaseous state alone.

Lattice Enthalpy

Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of solid ionic compound into the gaseous components. Lattice enthalpy of NaCl is -788 kJ mol-1 which means that 788 kJ of energy is required to separate 1 mole of NaCl into one mole of gaseous Na(g) and Cl-(g) to an infinite distance.

For Same Anion And Different Cations:

1. Cations having higher charge have large lattice energies than compounds having cations with lower charge. AlCl3 > CaCl2 > NaCl

2. As size of cation decrease, lattice energy increases. LiF > NaF > KF.

Teacher's Note

Lattice enthalpy is the energy needed to break apart one mole of an ionic solid. For example, to separate all the sodium and chloride ions in one mole of salt crystals requires 788 kJ of energy.

Exam Trick

Remember: Smaller cation + smaller anion = Larger lattice enthalpy = More stable compound. Think of it like strong tight packing of small things takes more energy to pull apart.

Points to Remember

Lattice enthalpy measures how much energy is needed to break an ionic solid into gaseous ions.
Smaller ions have higher lattice enthalpies because they are packed more tightly.
Ions with higher charges also have higher lattice enthalpies.
CsF has the weakest ionic bond despite high electronegativity difference.
Lattice enthalpy increases with decrease in size of ions.

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