Chapter 4: Chemical Bonding and Molecular Structure Short Questions | chemistry Class 11 CBSE

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Chapter 4: Chemical Bonding and Molecular Structure Short Questions | chemistry Class 11 CBSE | ByteNotes.in

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Question 1

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What is a chemical bond? Why do atoms form chemical bonds?

A chemical bond is the attractive force that holds various constituents (atoms, ions, etc.) together in different chemical species.
Atoms form bonds because every system tends to attain greater stability by lowering its energy; bonding is nature's way of achieving this lower energy state.
Under normal conditions, only noble gases exist as independent atoms; all other elements combine with atoms of the same or different elements to form molecules through chemical bonds.

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Question 2

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Write the Lewis dot symbols for Na, Mg, Al, N, O, and F.

Na has 1 valence electron (one dot), Mg has 2 valence electrons (two dots), Al has 3 valence electrons (three dots).
N has 5 valence electrons (three single dots and one pair), O has 6 valence electrons (two pairs and two single dots), F has 7 valence electrons (three pairs and one single dot).
The number of dots in each Lewis symbol represents the number of valence electrons, which determines the group valence of the element.

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Question 3

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State the octet rule and mention two of its limitations.

The octet rule states that atoms combine by transferring or sharing valence electrons in order to have eight electrons (an octet) in their valence shells, achieving noble gas configuration.
Limitation 1: Some compounds such as BF3 and AlCl3 have an incomplete octet around the central atom (fewer than eight electrons), violating the rule.
Limitation 2: In expanded octet molecules like PF5 and SF6, the central atom has more than eight electrons due to the availability of d orbitals in third-period elements.

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Question 4

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Define formal charge. Write the formula used to calculate it.

Formal charge is the hypothetical charge assigned to an atom in a polyatomic molecule or ion, assuming that bonding electron pairs are shared equally between atoms.
Formula: Formal Charge = (Total valence electrons of the free atom) - (Total non-bonding lone pair electrons) - (1/2)(Total bonding shared electrons).
Formal charges do not represent real charge separation; they are used to track valence electrons and to select the most stable Lewis structure (the one with smallest formal charges) from multiple possibilities.

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Question 5

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What is an electrovalent bond? Give one example with electron transfer equations.

An electrovalent (ionic) bond is formed as a result of the electrostatic attraction between a positively charged cation and a negatively charged anion produced by the transfer of electrons.
Example: Formation of NaCl — Na loses one electron to become Na+ (achieving Ne configuration), and Cl gains one electron to become Cl- (achieving Ar configuration).
The electrovalence is equal to the number of unit charges on the ion; Na has an electrovalence of +1 and Cl has an electrovalence of -1.

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Question 6

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Define lattice enthalpy. Why is it important for the stability of ionic compounds?

Lattice enthalpy of an ionic solid is defined as the energy required to completely separate one mole of the solid ionic compound into its gaseous constituent ions.
For NaCl, the lattice enthalpy is 788 kJ mol-1, meaning 788 kJ of energy is needed to separate one mole of solid NaCl into Na+(g) and Cl-(g) ions.
Even if the sum of ionization enthalpy and electron gain enthalpy is positive, ionic compounds are stable because the large energy released in lattice formation more than compensates for this; lattice enthalpy is thus the true measure of ionic compound stability.

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Question 7

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Distinguish between covalent radius and van der Waals radius.

Covalent radius is half the distance between two similar atoms joined by a covalent bond within the same molecule; it measures the size of the atom's core in a bonded state.
Van der Waals radius is half the distance between two similar atoms in separate adjacent molecules in a solid; it represents the overall size of the atom including its valence shell in a non-bonded situation.
For chlorine, the covalent radius is 99 pm while the van der Waals radius is 180 pm; van der Waals radius is always larger than covalent radius for the same element.

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Question 8

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Define bond enthalpy. How does it relate to bond strength?

Bond enthalpy is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in the gaseous state, expressed in kJ mol-1.
The H-H bond enthalpy in H2 is 435.8 kJ mol-1, the O=O bond enthalpy is 498 kJ mol-1, and the N triple bond N enthalpy is 946 kJ mol-1.
Larger bond dissociation enthalpy indicates a stronger bond; the relationship is direct — the stronger the bond, the greater the energy required to break it.

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Question 9

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State the relationship between bond order, bond enthalpy, and bond length.

Bond order is the number of covalent bonds between two atoms in a molecule; it is given by the number of shared electron pairs in the Lewis description.
As bond order increases, bond enthalpy increases — triple bonds are stronger than double bonds which are stronger than single bonds (e.g., C-C: 347, C=C: ~614, C triple bond C: ~839 kJ mol-1).
As bond order increases, bond length decreases — triple bonds are shorter than double bonds which are shorter than single bonds (e.g., C-C: 154 pm, C=C: 133 pm, C triple bond C: 120 pm).

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Question 10

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What is resonance? Why is a single Lewis structure inadequate for O3?

Resonance is the concept that when a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, and electron pairs (canonical forms) are written, and their hybrid represents the molecule accurately.
For O3, two Lewis structures can be drawn with one O-O single bond (148 pm) and one O=O double bond (121 pm) in each; but experimental data shows that both O-O bonds in O3 are equal at 128 pm, intermediate between single and double bond.
No single Lewis structure can represent this intermediate bond length, so the resonance hybrid of the two canonical forms describes O3 more accurately than either structure alone.

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Question 11

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Define dipole moment. Give its SI unit and express mathematically.

Dipole moment is the product of the magnitude of the charge (Q) and the distance of separation (r) between the centres of positive and negative charge in a polar molecule: mu = Q x r.
Dipole moment is a vector quantity, usually expressed in Debye units (D), where 1 D = 3.33564 x 10-30 C m (coulomb-meter in SI).
The dipole moment of H2O is 1.85 D (6.17 x 10-30 C m); non-polar molecules like H2, CO2, BF3, and CCl4 have zero dipole moment due to cancellation of individual bond dipoles.

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Question 12

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Why is the dipole moment of NF3 lower than that of NH3 despite F being more electronegative than H?

Both NH3 and NF3 are pyramidal molecules with a lone pair on nitrogen, but the orbital dipole due to the lone pair acts in opposite directions in the two molecules.
In NH3, the orbital dipole due to the lone pair acts in the same direction as the resultant N-H bond dipoles, making the net dipole moment large (4.90 x 10-30 C m).
In NF3, the orbital dipole due to the lone pair acts opposite to the resultant N-F bond dipoles, decreasing the net dipole moment (0.8 x 10-30 C m), giving NF3 a lower dipole moment than NH3.

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Question 13

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State the main postulates of the VSEPR theory.

The shape of a molecule depends on the number of valence shell electron pairs (bonded or nonbonded) around the central atom, and these pairs repel one another since their electron clouds are negatively charged.
Electron pairs tend to occupy positions in space that minimise repulsion and maximise the distance between them; the valence shell is taken as a sphere with electron pairs localised on its surface.
A multiple bond is treated as a single super pair for the purpose of geometry prediction; where resonance structures exist, VSEPR is applicable to any one of them.

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Chapter 4: Chemical Bonding and Molecular Structure | Short Questions

What is a chemical bond? Why do atoms form chemical bonds?

Write the Lewis dot symbols for Na, Mg, Al, N, O, and F.

State the octet rule and mention two of its limitations.

Define formal charge. Write the formula used to calculate it.

What is an electrovalent bond? Give one example with electron transfer equations.

Define lattice enthalpy. Why is it important for the stability of ionic compounds?

Distinguish between covalent radius and van der Waals radius.

Define bond enthalpy. How does it relate to bond strength?

State the relationship between bond order, bond enthalpy, and bond length.

What is resonance? Why is a single Lewis structure inadequate for O3?

Define dipole moment. Give its SI unit and express mathematically.

Why is the dipole moment of NF3 lower than that of NH3 despite F being more electronegative than H?

State the main postulates of the VSEPR theory.

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