Unit 3: Chemical Bonding — Short Questions

Ans. (i) Dipole
A molecule with δ+ charge on one part and δ- charge on other part is called dipole, e.g., HF, HCl, HBr

(ii) Bond order
Number of bonds formed between two atoms after the atomic orbitals overlaps, is called bond order. It is taken as half of the difference between the number of bonding electrons (say a) and antibonding electrons (say b)
Bond order = a - b / 2

(iii) Permanent Dipole-Permanent dipole force
The negative end of one permanent dipole attracts the positive end of other permanent dipole molecule. This force of attraction is called permanent dipole-permanent dipole force e.g.,
H—Cl·····H—Cl·····H—Cl
δ+ δ- δ+ δ- δ+ δ-

(iv) London dispersion force
Momentary force of attraction between instantaneous dipole molecule and induced dipole molecule is called London dispersion force. It is present in all type of molecules whether polar or non-polar.

Ans. (i) HCN (ii) NCℓ₃
H·:C≡N: :C̈ℓ:
:C̈ℓ:N:
:C̈ℓ:

(iii) CO (iv) O₃
:C̈≡Ö: ·Ö≡Ö—Ö:

(v) NO₂
:Ö:··N≡:Ö

Ans. (i) Xenon trioxide has a trigonal pyramidal shape due to lone pair-bond pair repulsion.

(ii)
δ+
/Xe
δ-O O δ-

Ans.
Sigma Bond:
• Sigma Bond is formed by head-to-head overlap of atomic orbitals along the internuclear axis.
• The electron density is maximum around the internuclear axis.

Pi Bond • Pi bond is formed by the lateral overlap of unhybridized p-orbitals.
• The electron density is located above and below the internuclear axis.

Ans. (i)
CH₃

C═O
/
CH₃
SiCℓ₄: SiCℓ₄ shows sp³ hybridization in which central silicon forms Four sigma bonds with four chlorin atoms as shown below.

(ii) Yes propanone can make hydrogen bond with water when both are intermixed as shown below
CH₃

C═O···········H—O
/
CH₃ H

Ans. SO₄²⁻
According to VSEPR Theory, the shape of SO₄²⁻ ion is tetrahedral. The central sulfur atom is surrounded by four oxygen atoms and there are no lone pairs on sulfur.

:O²⁻
|
:O²⁻—S—:O²⁻
|
:O²⁻
[shown as 2- charge indicator]

BH₄⁻
According to VSEPR Theory, the shape of ion is tetrahedral in which central Boron atom has four bond pair and no lone pair as shown below.
H
|
H—B—H
|
H
[shown as - charge indicator]

Tri-iodide ion (I₃⁻)
:I⁻:I:I:
[shown as - charge indicator]

Ans. Sideways overlap When the axes of two p-orbitals (i.e pᵧ or pz orbitals) are parallel to each other, they interact to form π molecular orbitals. The bonding molecular orbitals π(2pᵧ) or π(2pz) have zero electron density on the nuclear axis (called the nodal plane). On the other hand, anti-bonding molecular orbitals π*(2pᵧ) and π*(2pz) have the least electron density in the inter-nuclear region as in figure.

Ans. PCℓ₃
It shows sp³ hybridization because p has three sigma bonds and one lone pair as show below.

SiCℓ₄ SiCℓ₄ shows sp³ hybridization in which central silicon forms Four sigma bonds with four chlorin atoms as shown below.

[NH₄]⁺
NH₄ shows sp³ hybridization in which central Nitrogen atom forms four sigma bonds with four hydrogen atoms.
H
|
H—N—H [shown with + charge]
|
H

In above ions the central Nitrogen obey octet rule.

Ans. PF₃
δ+
/P
δ-F F δ-

F
δ-

PF₃ is a polar molecule its dipole moment is 0.985D

SiF₄
δ-
|
δ-—Si—δ-
|
δ-

SiF₄ is a non-polar molecule. It has zero dipole moment.

Ans. Shape = Linear Molecule
Angle = 180°

The CO₂ molecule shows sp hybridization where central C forms two sigma bonds and two pi bonds with oxygen atoms. The geometry is linear due to absence of lone pairs on central carbon.

Ans. (i) CℓO₄⁻
:O²⁻
|
:O—Cℓ—O:
|
:O
[shown with - charge]

In the above structure central Cℓ has expanded octets by eight electrons because Cℓ belongs to 3rd period

(ii) ICℓ₄⁻
In above structure I has expanded octet by four electrons because Iodine belongs to 5th period.
:Ċℓ:
:Cℓ—I—Cℓ:
:Ċℓ:
[shown with - charge]

(iii) NH₄⁺
H
|
H—N—H [shown with + charge]
|
H

In above ions the central Nitrogen obey octet rule.

(iv) I₃⁻
:I···I···I:
[shown with - charge]

Ans. The bond between K and Cℓ is ionic because their electron affinity difference is 2.34 which is greater than 1.8.
The bond between Si and Cℓ is polar because their electronegativity difference is 1.26 which fall between 0.4-1.8

Ans. SO₂ is polar due to its bent molecular geometry which results in a net dipole moment

Ö²⁻
|| δ+
Ö—S—O
δ+ δ+
P = 0.0 D

SO₃ is non-polar because of its trigonal planar shape allowing individual bond dipoles to cancel each other

Ö²⁻
/
δ+—S—δ+
/
Ö
δ+
P = 1.61 D

Ans. In the molecular orbital diagrams of O₂, and O₂⁻.
In the molecular orbital diagram of O₂, there are two unpaired electrons in π (2pᵧ) and π*(2pz) orbitals. So O₂ is paramagnetic. In the molecular orbital diagram of O₂, π(2pᵧ) and π*(2pz) orbitals are vacant. So paramagnetic behaviour vanishes and now it is diamagnetic. In case of O₂²⁻ two electrons are absorbed into O₂. Therefore π*(2pᵧ) and π*(2pz) will contain four electrons. There is no unpaired electron. So, it is not paramagnetic.

Ans.
Bone | Electronegativity Difference | Polarity
C-Cℓ | 0.66 | Slightly Polar
Si-F | 2.1 | Ionic
Se-F | 1.45 | Most Polar

Ans. Triple bond is stronger than double and single bound it how following order.
Triple bond > double bond > single bond
Single bond means one shared pair electron
Double bond means two shared pair electrons
Triple bond means three shared pair electrons
As the number of shared pair electrons increases it leads to stronger nuclear attraction and shorter bond lengths and higher bond energies so following trend observe.
Triple bond > Double bond > Single bond

Ans. According to the Lewis theory, the ionic bond is formed by the complete transfer of electrons from an, atom with low ionization energy to another atom with high electron affinity.

Ans. A dative bond is formed between two atoms when the shared pair of electrons is donated by one of the bonding atoms.

Ans. An ionic bond is formed by the complete transfer of one or more electrons from one atom to another, resulting in the formation of oppositely charged ions. This bond occurs between metals and nonmetals. For example, in NaCl, sodium donates one electron to chlorine, forming Na⁺ and Cl⁻ ions, which are held together by electrostatic attraction.

Ans. Ionic compounds consist of a regular lattice of alternating positive and negative ions. The strong electrostatic forces of attraction between oppositely charged ions require a large amount of energy to break. Therefore, ionic compounds are usually hard solids and have high melting and boiling points.

Ans. In solid form, ions in ionic compounds are fixed in position and cannot move freely. However, when melted or dissolved in water, the ions are free to move, allowing them to conduct electricity. The presence of mobile ions is essential for electrical conductivity in ionic substances.

Ans. Lattice energy is directly related to the strength of an ionic bond. The greater the lattice energy, the stronger the ionic bond. It depends on the size and charge of the ions; smaller and highly charged ions form lattices with higher lattice energies.

Ans. Covalent compounds do not contain free ions or electrons that can carry electric current. The atoms are held together by shared electron pairs in neutral molecules. Hence, in solid, liquid, or aqueous form, covalent compounds are poor conductors of electricity.

Ans. In a double bond, two electron pairs are shared between two atoms, while in a triple bond, three pairs are shared. Double and triple bonds are shorter and stronger than single bonds. For example, oxygen forms a double bond in O₂, and nitrogen forms a triple bond in N₂.

Ans. Bond order refers to the number of chemical bonds between a pair of atoms. A higher bond order means stronger and shorter bonds. For example, in N₂, the bond order is 3 (a triple bond), making it stronger than O₂, which has a bond order of 2.

Ans. Bond polarity arises when two atoms with different electronegativities share electrons unequally. The more electronegative atom attracts the shared electrons more strongly, developing a partial negative charge. For instance, in HCl, chlorine is more electronegative than hydrogen, making the bond polar.

Ans. A coordinate covalent bond is formed when one atom donates both of the electrons shared in a bond. In contrast, in regular covalent bonds, each atom contributes one electron. An example is the formation of NH₄⁺, where NH₃ donates a lone pair to H⁺.

Ans. One atom must have a lone pair of electrons (a donor), and the other must have an empty orbital (an acceptor). The donor provides both electrons to form a shared pair, resulting in a stable coordinate covalent bond.

Ans. Metallic bonding is the attraction between positively charged metal ions and a "sea" of delocalized electrons. Unlike ionic or covalent bonds, electrons are not shared or transferred between specific atoms. This delocalization gives metals their unique properties.

Ans. In metals, valence electrons are free to move throughout the structure. These mobile electrons carry electric current easily when a voltage is applied, making metals excellent conductors of electricity in solid and liquid states.

Ans. It is a quantitative measurement of the polarity of a bond or a molecule. The dipole moments of diatomic molecules like HF, HCl, HBr, HI, NO, etc. are directed from the positive ends (δ +) to negative ends (δ -). Dipole moment is measured in Debye unit (D). Higher dipole moment indicates high polarity in a molecule.

Ans. CCl₄ has four C-Cℓ bonds which are expected to have high polarity due to a large electronegativity difference between C and Cℓ atoms. The bond moments associated to the four C-Cℓ bonds are directed in such a way that they cancel each other. The net dipole moment of the CCl₄ molecule is zero making it a non-polar molecule. The CCℓ₄ molecule id perfectly tetrahedral.

Ans. BeCℓ₂ is a linear molecule having two similar Cℓ atoms on both sides of the central atom at 180°. The individual Be-Cℓ bond moments are cancelled out as they are opposite in direction and equal in amount. So, it has zero dipole moment.

Ans. BF3 has a trigonal plannar symetrical structure and its dipole moment is also zero. As a rule, the molecules that have same ligands (atoms or groups of atoms with the central zero) in a regular geometry, the individual dipole moments may not be zero, but overall molecule has zero dipole moment. Such a molecule is said to be non-polar.

Ans. Bond polarity arises when atoms with different electronegativities form a covalent bond. The more electronegative atom attracts electrons more strongly, creating a partial negative charge, while the other atom becomes partially positive. The greater the electronegativity difference, the more polar the bond.

Ans. Dipole moment is a measure of the polarity of a molecule. It is the product of the charge difference and the distance between the charges. A molecule with a higher dipole moment is more polar. It is represented by the Greek letter μ and measured in Debye units (D).

Ans. Water has a bent shape due to lone pairs on oxygen, so its bond dipoles do not cancel, making the molecule polar. In contrast, carbon dioxide is linear, and the dipoles of the two C=O bonds cancel each other out, resulting in a non-polar molecule despite having polar bonds.

Ans. From Lewis structures of SO₄²⁻, we can calculate the number of electrons around the central S atom. The number of electrons in the valence shell of S atom can be calculated as:
No. of valence electrons = 2 × (double bond electrons) + 2 × (single bond electrons)
= 2(4) + 2(2) = 12
Thus, S has 12 electrons in its valence shell and it exceeds the octet by 4 electrons.

Ans. In the same way, we can calculate the number of valence electrons around the central iodine atom in the tri-iodide ion. The Lewis structure of the ion is given by: The number of valence electrons = 2 × (no. of single bond electrons) + 2 × (no. of lone pair)
= 2(2) + 2(3) = 10
The central iodine atom in tri-iodide ion I3 has 10 electrons in the valence shell and the octet expands by two electrons

Ans. Definition

• Polyatomic ions are the ions composed of more than one type of atoms. There formal charge on them which is calculated based on the number of electrons in their valence shells after the formation of bonds. With the exception of ammonium ion (NH₄⁺), these ions mostly carry a negative charge, for example the carbonate (CO₃²⁻), sulfate (SO₄²⁻), and nitrate (NO₃⁻) ions.

Ans. The atoms having electronegativity difference less than 0.4 are said to make a pure covalent bond. An electronegativity difference between 0.4 and 1.8 corresponds to a polar covalent bond, whereas, above this value the bond between two atoms will be ionic in nature

Ans. Bond length is the distance between the nuclei of two bonded atoms. It decreases as the number of shared electron pairs increases. A single bond is longer than a double bond, which in turn is longer than a triple bond. For example, in carbon compounds: C-C > C=C > C≡C in bond length.

Ans. There is an inverse relationship between bond energy and bond length. A shorter bond is generally stronger and has a higher bond energy. As bond length increases, the attraction between nuclei and bonding electrons decreases, leading to a weaker bond and lower bond energy.

Ans. Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around the central atom repel each other and arrange themselves to be as far apart as possible. This determines the shape of the molecule to minimize repulsion.

Ans. Methane has a tetrahedral shape with bond angles of approximately 109.5°. This is because the four bonding pairs of electrons repel each other equally and arrange themselves as far apart as possible around the central carbon atom.

Ans. Water has two lone pairs and two bonding pairs on the central oxygen atom. Lone pairs repel more strongly than bonding pairs, distorting the shape into a bent geometry with a bond angle of about 104.5°.

Ans. Carbon dioxide (CO₂) has two double bonds on opposite sides of the carbon atom and no lone pairs on the central atom. The repulsion between the two electron regions results in a linear arrangement with a bond angle of 180°.

Ans. Hybridization is the mixing of atomic orbitals to form new orbitals (hybrids) that are suitable for bonding. It explains the shape and bond angles of molecules better than using unhybridized orbitals. For example, in CH₄, carbon uses sp³ hybridization to form four equivalent bonds.
σ bonds + lone pairs = 4

Ans. In NH₃, the nitrogen atom is sp³ hybridized because it has three sigma bonds and one lone pair. The molecule has a trigonal pyramidal geometry, and the bond angle is about 107°, slightly less than tetrahedral due to lone pair repulsion.

Ans. The electron configuration of He is 1s2. For a successful formation of He₂ molecule, 1s orbitals of two of He₂ molecule, 1s orbitals of two He atoms must combine to form bonding (σ 1s) and anti-bonding (σ* 1s) orbitals as shown in Figure 3.36. Out of four electrons, two enter the bonding molecular orbital σ (1s) and the remaining two occupy the antibonding δ*(1s) molecular orbital. But on calculation we discover that the bond order for He₂ is zero. Hence, He₂ molecule is not formed.
Bond order of He₂ = 2−2/2 = 0

Ans. The number of bonds formed between two atoms after the atomic orbitals overlap, is called the bond order and is taken as half of the difference between the number of bonding electrons (say a) and antibonding electrons (say b).
Bond Order = (No. of Electrons in B.M.O) - (No. of Electrons in A.B.M.O) / 2 = 2−0 / 2 = 1

Ans. The energy is the average amount of energy required to break all bonds of a particular type in one mole of a substance. It is determined experimentally by measuring the heat involved in a chemical reaction and its unit is KJ mol⁻¹. Bond energy is a measure of the strength of a bond and its reactivity
Bond length is the distance between the nuclei of two atoms forming a covalent bond.
The bond lengths are experimentally determined by physical techniques, such as electron diffraction, X-ray diffraction, or spectral studies. The bond length of a bond is governed by many factors including electronegativity, size, and the nature of the covalent bond (single, double, or triple).

Ans. In NF₃, the strong polarity of N-F bond pulls the lone pair of atom closer to its nucleus, which in turn exerts a stronger repulsion over bonding electrons. Thus, angle further shrinks to 102°. Moreover, the bond pairs N-F bonds are closer to F atom than N·atom. The increased distances in these bond pairs makes their repulsion less operative.

Ans. Chemical bond's, i.e. ionic, covalent, and metallic bond usually termed as true chemical bonds, as the affect the chemical properties of a substance. This is because of certain reasons. Firstly, chemical bonds result in the formation a new species through transfer and sharing of electrons. Whereas, intermolecular forces act to bring molecules closer and influence the physical properties. However, such a clear distinction between the chemical bonds and intermolecular forces is not possible.

Ans. See Q.8 of theory.

Ans. See Q.13,14,15 of theory.

Ans. See Q.9,10,11 of theory.

Ans. See Q.18 of theory.

Ans. See Q.19 of theory.