Chapter 4: Chemical Bonding and Molecular Structure

Class 11 Chemistry Chapter 4: Chemical Bonding and Molecular Structure Internal Problems and Answers : 

Problem 4.1 : Write the Lewis dot structure of CO molecule.

Problem 4.2 : Write the Lewis structure of the nitrite ion,   .

problem 4.3 :  Explain the structure of  ion in terms of resonance.

problem 4.4 : Explain the structure of  molecule.

Class 11 Chemistry Chapter 4: Chemical Bonding and Molecular Structure Exercise Questions and Answers : 

4.1 Explain the formation of a chemical bond.

Answer: Chemical bonds form when atoms interact to achieve stable electron configurations, often resembling noble gas arrangements. Atoms form bonds by sharing, donating, or accepting electrons to lower their energy. Ionic bonds occur when electrons transfer from one atom to another, producing oppositely charged ions that attract. Covalent bonds arise when atoms share electrons, resulting in polar or nonpolar bonds based on electron distribution. Metallic bonds involve a ‘sea’ of delocalized electrons moving freely among metal ions. This process reduces overall system energy significantly and effectively.

4.2 Write Lewis dot symbols for atoms of the following elements : Mg, Na, B, O, N, Br.

Answer:

4.3 Write Lewis symbols for the following atoms and ions:  S and  ; Al and  ; H and

4.4 Draw the Lewis structures for the following molecules and ions :

4.5 Define octet rule. Write its significance and limitations.

Answer:  Atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing valence electrons in order to achieve eight electrons in their valence shells. This tendency of atoms to attain a stable octet configuration is known as the octet rule.

Significance of the Octet Rule:

Explains Bond Formation: Helps in understanding how atoms form ionic and covalent bonds.

Predicts Stability: Atoms with a complete octet tend to be more chemically stable.

Determines Molecular Structure: Aids in predicting the shape and structure of molecules.

Limitations of the Octet Rule:

Incomplete Octet: Elements like H (2e⁻), Be (4e⁻), B (6e⁻) do not always complete an octet.

Expanded Octet: Elements beyond period 2 (e.g., PCl₅, SF₆) can have more than eight valence electrons.

Odd-Electron Molecules: Some molecules like NO, NO₂ contain an odd number of electrons, violating the rule.

Transition Metals Exception: Many transition metals do not follow the octet rule due to their complex electron configurations.

4.6 Write the favourable factors for the formation of ionic bond.

Answer: The favorable factors for the formation of an ionic bond are :

(i) Large Electronegativity Difference: A significant difference in electronegativity between the atoms (e.g., metal and non-metal) encourages electron transfer.

(ii) Low Ionization Energy: The metal atom easily loses electrons to form cations.

(iii) High Electron Affinity: The non-metal gains electrons readily to form anions.

(iv) Small Size of Cations and Anions: Smaller ions allow closer packing, stabilizing the ionic bond.

(v) High Lattice Energy: A strong electrostatic attraction between oppositely charged ions enhances bond formation.

4.7 Discuss the shape of the following molecules using the VSEPR model:  

4.8 Although geometries of  and  molecules are distorted tetrahedral, bond angle in water is less than that of ammonia. Discuss.

4.9 How do you express the bond strength in terms of bond order ?

Answer: Bond strength can be expressed in terms of bond order, which is defined as the difference between the number of bonding electrons and antibonding electrons, divided by two. A higher bond order indicates a stronger bond because more electrons are involved in bonding and fewer in antibonding orbitals. For example, a bond order of 1 corresponds to a single bond (weaker), 2 corresponds to a double bond (stronger), and 3 to a triple bond (strongest).

4.10 Define the bond length.

Answer: Bond Length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. It is the distance at which the attractive and repulsive forces between the atoms are balanced, resulting in a stable bond. Bond lengths are typically measured using spectroscopic, X-ray diffraction, and electron-diffraction techniques. Each atom contributes to the bond length, with the covalent radius of each atom representing its contribution in the case of a covalent bond.

4.11 Explain the important aspects of resonance with reference to the  ion.

4.12  can be represented by structures 1 and 2 shown below. Can these two structures be taken as the canonical forms of the resonance hybrid representing  ? If not, give reasons for the same.

   

4.13 Write the resonance structures for   and   .

4.14 Use Lewis symbols to show electron transfer between the following atoms to form cations and anions : (a) K and S (b) Ca and O (c) Al and N.

4.15 Although both   and  are triatomic molecules, the shape of  molecule is bent while that of  is linear. Explain this on the basis of dipole moment.

4.16 Write the significance/applications of dipole moment.

Answer: The dipole moment is a measure of the separation of positive and negative charges in a molecule, indicating its polarity.

Its significance and applications are :

Polarity of Molecules: It helps determine whether a molecule is polar or nonpolar, influencing its chemical behavior and interactions.

Intermolecular Interactions: Affects the strength of intermolecular forces (e.g., dipole-dipole, hydrogen bonding) in liquids and solids.

Solubility: Polar molecules tend to dissolve in polar solvents (like water), while nonpolar molecules dissolve in nonpolar solvents.

Chemical Reactivity: The dipole moment can influence a molecule's reactivity in chemical reactions, especially in polar environments.

Spectroscopic Analysis: It plays a crucial role in techniques like microwave spectroscopy for identifying molecular structure.

4.17 Define electronegativity. How does it differ from electron gain enthalpy ?

Answer: Electronegativity is the tendency of an atom in a molecule to attract shared electrons towards itself.

Differences between Electronegativity and Electron Gain Enthalpy:

Property	Electronegativity	Electron Gain Enthalpy
Definition	Ability of an atom to attract shared electrons in a bond.	Enthalpy change when an electron is added to a neutral atom.
Units	No units (dimensionless).	Measured in kJ/mol.
Property Type	Relative measure of attraction in a bond.	Quantitative measure of energy change.
Dependence	Depends on nuclear charge, size, and electron shielding.	Depends on the atom's ability to accept an electron.
Example	Fluorine has the highest electronegativity (3.98 on the Pauling scale).	Chlorine has a highly negative electron gain enthalpy (-349 kJ/mol).

 

4.18 Explain with the help of suitable example polar covalent bond.

Answer:

4.19 Arrange the bonds in order of increasing ionic character in the molecules:  ,  ,  ,  and  .

4.20 The skeletal structure of  as shown below is correct, but some of the bonds are shown incorrectly. Write the correct Lewis structure for acetic acid.

  

4.21 Apart from tetrahedral geometry, another possible geometry for  is square planar with the four H atoms at the corners of the square and the C atom at its centre. Explain why  is not square planar ?

4.22 Explain why  molecule has a zero dipole moment although the Be–H bonds are polar.

4.23 Which out of  and  has higher dipole moment and why ?

4.24 What is meant by hybridisation of atomic orbitals? Describe the shapes of  hybrid orbitals.

4.25 Describe the change in hybridisation (if any) of the Al atom in the following reaction. 

  

4.26 Is there any change in the hybridisation of B and N atoms as a result of the following reaction?

  

4.27 Draw diagrams showing the formation of a double bond and a triple bond between carbon atoms in  and  molecules.

4.28 What is the total number of sigma and pi bonds in the following molecules? (a)   (b)  

4.29 Considering x-axis as the internuclear axis which out of the following will not form a sigma bond and why? (a)  and  (b)  and   ; (c)    and  

(d) 1s and 2s.

4.30 Which hybrid orbitals are used by carbon atoms in the following molecules?

(a)     (b)     (c)     (d)   (e)

4.31 What do you understand by bond pairs and lone pairs of electrons? Illustrate by giving one exmaple of each type.

4.32 Distinguish between a sigma and a pi bond.

4.33 Explain the formation of  molecule on the basis of valence bond theory.

4.34 Write the important conditions required for the linear combination of atomic orbitals to form molecular orbitals.

4.35 Use molecular orbital theory to explain why the molecule does not exist.

4.36 Compare the relative stability of the following species and indicate their magnetic properties :   (superoxide), (peroxide)

4.37 Write the significance of a plus and a minus sign shown in representing the orbitals.

4.38 Describe the hybridisation in case of  . Why are the axial bonds longer as compared to equatorial bonds?

4.39 Define hydrogen bond. Is it weaker or stronger than the van der Waals forces?

4.40 What is meant by the term bond order? Calculate the bond order of :   and   .