Unit 3: Chemical Bonding — Long Questions

Ans. Chemical Bond
Definition: A chemical bond is the force that holds together two or more atoms, molecules or ions.
The property is of a substance depend on the type of the chemical bond between its atoms.
Explanation: The term chemical bond includes intra-molecular forces which are:
(i) Ionic bond
(ii) Covalent bond
(iii) Dative bond
(iv) Metallic bond

TYPES OF BONDING

Lewis concept
• Lewis concept of bonding gives a simple explanation of the formation of all types of bonds.
• Atoms make bonds to complete their outermost shells to have noble gas-like configuration.
• It is mostly attained through the formation of an octet in the valence shell.

1. Ionic Bond (electrovalent bond)
Definition: 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.
Example: Formation of ionic bond in NaCℓ
• The Na atom (₁₁Na = 2,8,1) tends to lose the outermost electron to form Na⁺ (2,8) ion, which has the electron configuration of ₁₀Ne, a noble gas nearest to it.
• Chlorine atom Cℓ (2,8,7) gains one electron to form the chloride ion, Cℓ⁻ (2,8,8), also gaining the next noble gas electron arrangement.
• The oppositely charged Na⁺ and Cℓ⁻ ions are held together by strong ionic bond in the crystal of NaCℓ.
The formation of ionic bond is shown in the figure.
A similar type of bond is expected between group 1 and 2 metals and groups 17 and 16 non-metal elements.

2. Covalent Bond
Definition: A covalent bond is formed by the mutual sharing of electrons between two atoms.
By sharing electrons, each atom completes its valence shell and attains the nearest noble gas configuration.
Types:
(i) Polar Covalent bond
(ii) Non-Polar Coveant bond

(i) Non Polar bond
Definition: The bond formed due to equal sharing of electrons between two atoms.
• In formation of Cℓ₂ molecule, the bond formation between two Cℓ atoms is shown in figure.
• The bond between two Cℓ atoms purely covalent and non-polar.
• The electronegativity of the two atoms is exactly the same.
• So, the bonded atoms remain electrically neutral and there is no charge on either atom.
The other such molecules are H₂, F₂, Br₂, and I₂.

(ii) Polar Covalent bond
Definition: The bond formed due to unequal sharing of electrons between two atoms.
Examples of polar covalent bonds are H-Cℓ, H-Br, H₂O, etc.

3. Dative Bond (Co-ordinate Covalent Bond)
Definition: A dative bond is formed between two atoms when the shared pair of electrons is donated by one of the bonding atom only.
Examples of dative bond: Following are the examples of the dative bond.

(i) Hydronium ion (H₃O⁺)
• In the co-ordinate covalent bond formation between H₂O and a proton (H⁺). H₂O has two covalent bonds and there are two lone pairs of electrons on the oxygen atom.
• The proton is deficient in electrons.
• The oxygen atom can donate its lone pair of electrons to the acceptor H⁺.
• This results in the formation of a dative bond as in the figure.

(ii) Carbon monoxide
• Carbon monoxide also contains a dative bond between oxygen and carbon atoms.
• Oxygen shares its two valence shell electrons with the carbon atom to make a normal double covalent bond.
• The valency of carbon is not satisfied with these bonds.
• It needs one more pair of electrons to complete its octet.
• Oxygen atom donates one of its lone pairs for the formation of a covalent bond, which is a dative bond as exhibited in figure.

(iii)Ozone
• It is formed by a reaction between the oxygen molecule with an oxygen atom (a free radical (O•).
• One of the Oxygen atoms in the molecule donates a pair of electrons to make a coordinate covalent bond with this free radical.
• The central atom has total three bonds including a double covalent bond on one side and a coordinate covalent bond on the other.
• It carries a positive charge, while the oxygen atom that accepts the lone pair carries a negative charge.
• The structure of ozone molecule is presented in figure.

Ans. Polyatomic Ions
Definition: Polyatomic ions are the ions composed of more than one type of atoms. There formal charge is the net charge on them which is calculated based on the number of electrons in their valence shells after the formation of bonds. Except ammonium ion (NH₄⁺), these ions mostly carry a negative charge.
Example: The examples of polyatomic ions are carbonate (CO₃²⁻), sulfate (SO₄²⁻), and nitrate (NO₃²⁻) ions.

Expanded Octets
• In some polyatomic ions, the central atom violates the octet rule by expanding its electrons density to the higher orbitals.
• These are set to have expanded octets.

Examples of expanded octets The prominent examples of expanded of octets are SO₄²⁻, C/O₄⁴⁻, PO₄³⁻

(i) Expanded Octets in SO₄ 2-:
From Lewis structures of SO₄²⁻, the number of electrons around the central S atom can be calculated.
The number of electrons in the valence shell of S atom can be no. of calcuated as:
No. of valence electrons = 2 × (double bond electrons) + 2 × (no. of single bond electrons)
= 2(4) + 2(2) = 12 electrons
• S has 12 electrons in its valence shell and it exceeds the octet by 4 electrons.
• The expansion of octet is caused by the involvement of the d orbital in bonding, which can accommodate the extra electrons.
• The following electronic configuration of native S (0) atom shows that it has two unpaired electrons in the p orbitals.
• The presence of d orbital allows this configuration to extend to 4 unpaired electrons by the transfer of one electron from 3s and one from the 3p electron pair.

It explains not only the variable oxidation states of S, but also the possibility of accepting extra electrons in its d orbital.

(ii) Expanded Octets in I₃⁻¹ Ion:
The number of valence electrons around the central iodine atom in the tri-iodide ion can be calculated. The Lewis structure of the ion is given by:
The number of valence electrons of the central atom can be calculated as follows.
No of valence electrons = 2 × (no. of single bond electrons) + 2 × (no. of lone pair)
= 2(2) + 2(3) =10 electrons
The central iodine atom in tri-iodide ion I₃⁻ has 10 electrons in the valence shell and the octet expands by two electrons.

Ans. Electronegativity and the type of bonds
The difference in electronegativity of two bonded atoms provides an approximate measure of the bond polarity and an indication of the type of bond.
Explanation
• When this difference is very small or zero, the bond is covalent and non-polar.
• When it is large, the bond is polar covalent or ionic. The following figure graphically depicts the change in the nature of the bond with electronegativity difference.
• 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.
• The electronegativity differences between the atoms in the bonds such as H-H bond is purely non polar.
• H-Cℓ a polar covalent bond, whereas NaCℓ is an ionic compound.

Ans. Dipole Moment and polarity of molecules
Definition: It is a quantitative measurement of the polarity of a bond or a molecule. A molecule with δ+ charge on one part and δ – charge on the other part is called a dipole and such a molecule is said to have a dipole moment.

Direction of dipole moment
The dipole moments of diatomic molecules like HF, HCl, HBr, HI, NO, etc. are directed from the positive ends (δ +) to negative ends (δ -).

Unit Dipole moment is measured in Debye unit (D). Higher dipole moment indicates high polarity in a molecule.

Dipole Moment and structure of molecules
Some molecules have net zero dipole moment as the symmetry in the structure causes the cancelation of the individuals bonds moments.
For compounds having more than two atoms, the dipole moment depends on the structure of the molecule.

Dipole Moment of H₂O: The dipole moment of water is 1.85 D which is directed from the end having two hydrogen atoms to the end with the oxygen atom as in the structure.
• A linear H₂O molecule (H-O-H) would have zero dipole moment.
• The non-zero dipole moment value shows that water is a non-linear molecule.
• Experiments reveal that the water molecule has a V-shaped structure.

Dipole Moment of SO₂ In contrast, SO₂, has a dipole moment of 1.61 D and in opposite direction to that in water.

Dipole moment of H₂S: HS is also a non-linear molecule as the individual bonds are polar, but they don't cancel each other's dipole moment and the overall molecule has some dipole moment.

Dipole Moment of BeCℓ₂
• 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.

Dipole Moment of CCℓ₄
• 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 is perfectly tetrahedral.

Dipole Moment of BF₃
• BF₃ has a trigonal plannar symmetrical 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.
• The overall molecule has zero dipole moment.
• This molecule is said to be non-polar.

Ans. Intermolecular forces
Definition: The forces of attraction between two molecules in a compounds are called intermolecular forces.

Explanation
Intermolecular forces are also considered as binding forces just like chemical bonds.
• The attraction between the molecules is much weaker than the chemical bonds.
• These forces are believed to exist between all kinds of atoms and molecules when they are sufficiently close to each other.
• Such intermolecular forces are called van der Waals forces.
• These intermolecular forces bring the molecules close together and give particular physical properties to the substances in gaseous, liquid, and solid states.

Three types of such forces are mentioned here:
(i) Instantaneous dipole-induced dipole forces or London dispersion forces
(ii) Permanent dipole-permanent dipole forces
(iii) Hydrogen Bonding

Ans. Bond energy
Definition: The Bond energy is the average amount of energy required to break all bonds of a particular type in one mole of a substance.
• Measurement of bond energy: 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.
• The strength of a bond depends upon the following factors:
(a) Electronegativity difference of bonded atoms
(b) Size of atoms

(a) Electronegativity difference of bonded atoms
• The bond energies of H-X type of compounds can be observed, where X= F, Cℓ, Br, I.
• The data shows that the bond energy of a bond rises with the increasing electronegativity difference between the bonded atoms.
• As the difference in electronegativity increases, the bond polarity also becomes greater and this gives rise to additional attractive force for binding the atoms. This is why the bond energy for HF is higher (565 kJ mol⁻¹) than for HI (295 kJ mol⁻¹).

Multiple bonds It may be noted that energies of multiple bonds are greater than those of single bonds.
C≡C > C=C > C-C

Bond Length
Definition: Bond length is the distance between the nuclei of two atoms forming a covalent bond.

Measurement of bond length The bond lengths are experimentally determined by physical techniques, such as electron diffraction, X-ray diffraction, or spectral studies.

Factors 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). Some selected bond lengths are given in table 3.2.

Electro-negativity cause shortening of bond.
With an increase in size of the atoms, the covalent bond length also increases.
The C-Cℓ bond length is about 180 pm, whereas the C-F bond length is 135 pm. This is because Cℓ atom is much larger.

Ionic character of Si-F bond: The bond length of Si-F bond in SiF₄, is found to be near 155pm, whereas the calculation of Si-F bond forms the covalent radii of Si and F (Si-117 pm and F-64 pm) is 181 pm.

The electronegativity difference causes an ionic character in the covalent bond.
The ionic character results in shortening of the bond length due to the additional attraction between the bonded atoms.

Ans. Chemical bond
Chemical bond's, i.e. ionic, covalent, and metallic bond usually termed as true chemical bonds, as they affect the chemical properties of a substance. This is because of certain reasons.
Chemical bonds result in the formation a new species through transfer and sharing of electrons.

Intermolecular forces
Intermolecular forces act to .bring molecules closer and influence the physical properties. A clear distinction between the chemical bonds and intermolecular forces is not possible.

Comparison of chemical bonds and intermolecular forces
The distinctive feature of the chemical bonds and intermolecular forces is their bond strength. The strength of a force is measured by the bond energy. Table 3.3 provides a comparison among different chemical bonds and intermolecular forces.

Bond energy comparison

Ionic bond The bond energy of ionic bond (sodium chloride, 760 kJ mol⁻¹) is highest, followed by covalent bond (O-H=464 kJ mol⁻¹), and average hydrogen bond energy (20-25 kJ mol⁻¹). It shows that the ionic bond is the strongest form of chemical bonding. It also reveals that chemical bonds are generally much stronger than intermolecular forces.

Intermolecular forces
• The permanent dipole-dipole forces (5-20 kJ mol⁻¹) and London dispersion forces (1-20 kJ mol⁻¹) are even weaker as indicated by their low bond energies.
• The metallic bond is mostly elaborated in terms of electrostatic forces, although some theories suggest that it may be of covalent nature.
• In any case, the metallic bond is weaker than both the ionic and covalent bonds. The average bond energy of the metallic bond is 100-150 kJ mol⁻¹. A satisfactory argument for the low strength of metallic pool⁻¹. A satisfactory argument for the low strength of metallic bond is extensive delocalization of electrons within the metallic crystal.

Ans. Introduction
• The VSEPR model predicts and explain the shapes of molecules but does not reasons for the formation of bonds.
• VBT is concerned with both the formations of bonds and the shapes of molecules.
• This method describing a covalent bond considers a molecule as a combination of atoms. The postulates of VBT are given below:

Postulates • Formation of Covalent Bond: A covalent bond is formed when half-filled orbitals in the valence shell of two atoms with similar energy overlap.
• Strength of Bond: A greater overlap of the orbitals results in a stronger bond.
• Direction of Covalent Bond: Covalent bonds are directional. The direction of the bond is determined by the shape and mode of the overlapping orbitals.

Explanation of VBT

Formation of a sigma bond: A sigma bond is formed by the overlap of two half-filled atomic orbitals on adjacent atoms. The orbitals approach each other on the nuclear axis. Both s and p orbitals can overlap head on to form sigma bonds.

s-s Overlap: The s-orbital of the one atom overlap with the s orbital of the other to give a sigma (σ) bond.

(i) Hydrogen (H₂): This type of overlap occurs during the formation of H₂ molecule, where each hydrogen atom has a half filled 's' orbital. After bond formation electrons are paired up and the electron density in this molecule is symmetrical around the two nuclei as shown in figure.

s-p Overlap:
The s-orbital of the one atom overlap with the p orbital of the other to form sigma bond.

(ii) Hydrogen chloride (HCℓ): In HCℓ molecule, a half filled 's' orbital of hydrogen overlaps with the half-filled p orbital of chlorine as shown in figure. The electron density is higher close to Cℓ atom due to its higher electronegativity value. This is why HCℓ is a polar molecule.

p – p Overlap: The p-orbital of the one atom overlap with the p orbital of the other atom to form sigma bond.

(iii)Chlorine (Cℓ₂): An example of this type of overlap is the formation of the Cℓ₂ molecule where 'p' orbitals of two chlorine atoms overlap on the nuclear axis. The electron density is symmetrical around the nuclei of the two Cℓ atoms because both have same electronegativity value.

Interesting information!
S.Q. Why these is difference in bond angles of H₂O and H₂S?
Ans. The bond angle in H₂O is 104.5°, whereas in H₂S, it is 92°. This is because the orbitals of S are larger and the lone pair exerts a stronger repulsion.

(iv)H₂S
The H₂S molecule is a non-linear molecule which is formed by the combination of one sulfur and two hydrogen atoms. The two 3p (say 3pᵧ and 3pᵤ) orbitals of sulfur containing one electron each can overlap with the 1s orbitals of two hydrogen atoms. A v-shaped molecule is thus formed having a bond angle of 92° as in figure.

(v) O₂
Formation of π bond
• In the formation of a double covalent bond between oxygen atoms (O=O).
• There are two unpaired electrons on each atom in perpendicular p orbitals (say Pₓ and Pᵧ).
• The pₓ orbitals on the two oxygen atoms are oriented in such a way that they overlap end-to-end linearly.
• The linear overlap of the pₓ orbitals gives a σ bond.
• The pᵧ orbitals on both atoms are aligned parallel to each other.
• They overlap in a parallel way so that the two p lobes overlap above and the plane of the nuclei and the other two lobes below the plane as shown in figure.
• This results in the formation of π bond.
• The oxygen atoms are doubly bounded through σ and π bounds.

(vi) Hydrogen fluoride (HF): In the formation of H-F, the partially filled 1s orbital of H-atom overlap with partially filled orbital of fluorine to form sigma bond

(vii) Nitrogen (N₂): In case of N₂ p orbitals are half filled. One half-filled 2pₓ orbitals of one nitrogen overlap linearly with half-filled 2pₓ orbitals of second nitrogen. There are also two half-filled 2pᵧ and 2pᵤ orbitals which undergo sideways overlap to from two π bonds.

Ans. Hybridization
Definition: A process in which atomic orbitals of slightly different energies and shapes are mixed together to form a new set of equivalent orbitals of same energy and same shape is called hybridization.

Keep in mind!
S.Q. Give basic rules of hybridization.
Ans.
• Hybridization is a process of mixing of orbitals in a single atom (or ion).
• Only orbitals of comparable (relatively close) energies can be mixed to form hybrid orbitals.
• The number of mixing orbitals is always equal to the number of the resulting hybrid orbitals.

Types of Hybridization The types of hybridization are sp³, sp² and sp.

sp Hybridization
Definition: In sp hybridization, one s (low energy and spherical) and one p orbital (high energy and dumb-bell shaped), intermix to give a new set of two hybrid orbitals of same energy and same shape called sp hybrid orbitals.
These sp hybrid orbitals are arranged in linear geometry and oriented at 180°.
Examples: BeCℓ₂, CH = CH, MgCℓ₂ etc.

BeCℓ₂ In beryllium chloride the central atom is beryllium. The ground, excited and hybridized states of beryllium are given below.

The two sp hybrid orbitals lie linearly as in the following diagram.
BeCℓ₂ is formed when two sp hybrid orbitals of Be atom overlap with the half-filled p-orbitals of chlorine atoms.
The sp hybridization explains the geometry of linear molecules such as beryllium chloride, BeCℓ₂.
BeCℓ₂ is formed when two sp hybrid orbitals of Be atom overlap with the half-filled p-orbitals of chlorine atoms.

Q10. Explain BF₃ molecule with reference to sp²-hybridization.
Ans. sp² Hybridization
Definition: In sp² hybridization, one s and two p-orbitals of an atom intermix to form three orbitals called sp² hybrid orbitals. It is called sp² hybridization.
The three half-filled sp² hybrid orbitals are arranged in a trigonal planar geometry with bond angles of 120°.
Examples: BF₃, CH₂ = CH₂ BH₃, BC₍₃₎, propanone etc.

BF₃ The sp² hybridization explains the geometry of planar molecules such as BF₃. Electric configuration of sB.

The three outermost orbitals of B mix together to give three sp² hybrid orbitals. On the other hand, one of the p orbitals of fluorine is half filled (2p) orbitals.
BF₃ is formed by the overlap of three half-filled sp² hybrid orbitals of boron with 2pₓ, 2pᵧ orbitals of three fluorine atoms.

Note Answer continues beyond page 17.

Ans. sp² Hybridization
Definition: In sp² hybridization, one s and two p-orbitals of an atom intermix to form three orbitals called sp² hybrid orbitals. It is called sp² hybridization.
The three half-filled sp² hybrid orbitals are arranged in a trigonal planar geometry with bond angles of 120°.
Examples: BF₃, CH₂ = CH₂ BH₃, BCl₃, propanone etc.

BF₃ The sp² hybridization explains the geometry of planar molecules such as BF₃. Electric configuration of ₅B.

₅Bground = |↑↓| |↑| | | | Ground
2s 2p_x 2p_y 2p_z

Bexcited = |↑| |↑| |↑| | Excited Bhybridized = |↑| |↑| |↑| hybridized
2s 2p_x 2p_y 2p_z sp² sp² sp²

The three outermost orbitals of B mix together to give three sp² hybrid orbitals.
On the other hand, one of the p orbitals of fluorine is half filled (2p) orbitals.
BF₃ is formed by the overlap of three half-filled sp² hybrid orbitals of boron with 2p_x, orbitals of three fluorine atoms.

The structure is triangular planar with bond angles equal to 120°, each.

Ans. sp³ Hybridization
Definition: In sp³ hybridization, one s-orbital (low energy and spherical) and three p-orbitals (high energy and dumbbell shaped), intermix to give a new set of four orbitals of same energy and same shape called sp³ hybrid orbitals.
Each of these hybrid orbitals is a sp³ hybrid orbital. These hybrid orbitals are arranged in the tetrahedral geometry with angle of 109.5°.
Example: CH₄, NH₃, H₂O, NH₄⁺ etc. are compounds which show sp³ mode of hybridization.

Methane (CH₄)
The central atom is carbon. The electronic configuration of valence shell of ₆C in its excited and hybridized states are given as follow.

⁶Cground= |↑↓| |↑| |↑| | Ground
2s 2p_x 2p_y 2p_z

Cexcited = |↑| |↑| |↑| |↑| Excited Chybridized = |↑| |↑| |↑| |↑| hybridized
2s 2p_x 2p_y 2p_z sp³ sp³ sp³ sp³

Each of these hybrid orbitals is a sp³ hybrid orbital and they are arranged in the tetrahedral geometry.
The energies of hybrid orbitals are lower than unhybrid orbitals. The outermost four atomic orbitals of carbon mix up to give four hybrid orbitals of same energy and same shape.

The four new hybrid orbitals of equal energy have a tetrahedral geometry with carbon nucleus at the center. The methane (CH₄) molecule is formed by the overlap of sp³ hybrid orbitals of carbon with 1s orbitals of four hydrogen atoms separately to form four sigma bonds. The four C-H bonds, which result sp³ -s overlaps, are directed towards the corners of a regular tetrahedon.

Did you Know!
S.Q. What are the lobes of sp³ hybrid orbitals?
Ans. Each sp³ hybrid orbital consists of two lobes, one larger and the other smaller. For the sake of simplicity, the smaller lobe is usually not shown while representing sp³ hybrid orbitals together.

Ans. VSEPR Theory
Introduction: The VSEPR model describes the shapes of molecules based on the electron pairs that surround the central atom.
This model was presented by Sidgwick and Powell in 1940.
Assumptions: It is based on the assumption that a molecule will take a shape such that the electronic repulsions among the valence electrons of that atom are minimum.
In order to have the minimal repulsions, the electron pairs arrange themselves at farthest possible distances.
This arrangement of electron pairs determines the geometry of the resulting molecule.

Postulates
(i) Both the lone pairs (non-bonded) as well as the bond pairs participate in determining the geometries of molecules.
(ii) The electron pairs are arranged around the central atom so as to remain at maximum distance apart to avoid repulsions.
(iii) The electrons of lone pairs occupy more space than the bond pairs. As a result, the non- bonding electron pairs exert greater repulsive forces on bonding electron pairs and, thus, tend to compress the bond pairs. The magnitude of repulsions between the electron pairs in a given molecule decreases in the following order:
Repulsion Order: Lone pair-lone pair (lp-lp) > lone pair – bond pair (lp-bp) > bond pair – bond pair (bp-bp).

An electron pair shared by two nuclei occupies less space than a lone pair bound by a single nucleus.

(iv) The two electron pairs of a double bond and three electron pairs of a triple bond, contain a higher electronic charge density, but they are regarded as single pairs.
(v) The presence of highly electronegative atoms with the central atom results in decreased electronic repulsions between the bond pairs, but stronger repulsion by the lone pair.

Predicting the shapes of molecules
• In order to illustrate this model, the central atom is named 'A'. The electron pairs around 'A' are designated as 'B'.
• There may be different number of electron pairs around the central atom depending upon its valency.
• It gives rise to various types of molecules, such as AB₂, AB₃, AB₄, etc.
• The electron pair geometry of a molecule is determined by the total number of electron pairs around 'A'.
• The actual shape of the molecule is determined by the atoms excluding the lone pairs.
Following table 3.4 gives the possible shapes of different types of molecules having varying numbers of bond and lone pairs.

Table 3.4 Geometries and shapes of different systems of covalently bonded compounds

No. of electron pairs around central atom | No. of lone pairs | Electron pair geometry | Shape with respect to atoms | Angular separation of atoms | Examples
2 | 0 | Linear | Linear | 180° | CO₂, BeCl₂
3 | 0 | Trigonal | Trigonal | 120° | BCl₃, SO₃
3 | 1 | Trigonal | Angular, v-shaped | Less than 120° | SO₂, SnCl₂
4 | 0 | Tetrahedral | Tetrahedral | 109.5° | CH₄, CCl₄
4 | 1 | Tetrahedral | Pyramidal | Less than 109.5° | NH₃, H₃O⁺
4 | 2 | Tetrahedral | Angular, v-shaped | Less than 109.5° | H₂O, OF₂
5 | 0 | Trigonal bipyramidal | Trigonal bipyramidal | 120°, 90° | PCl₅, I₃⁻
6 | 0 | Octahedral | Octahedral | 90° | SF₆, XeF₄

Steps to Determine the shapes of molecules.
The following steps are followed to predict the shape of a molecule.
• The least electronegative atom or the element with the least number of atoms is mostly selected as the central atom. In CCl₄, the carbon atom is the central atom.
• The total number of electron pairs around the central atom are counted (including bond pairs and lone pairs). In CCl₄, the total number of bond pairs around the C atom are 4 and it has no lone pair.
• The total number of electron pairs determines the electron pair geometry (as in the above table) of the molecule.
• CCl₄ has total four electron pairs and its geometry is tetrahedral.
• The actual shape of the molecule is determined excluding the lone pairs (if any).

Ans. AB₂ type molecules (linear geometry)
In such molecules, two electron pairs around the central atom are arranged at an angle of 180° to minimize repulsions between them. Thus, they form a linear geometry as shown in figure.
Example: BeCl₂ molecule is of AB₂ type with Be as the central atom and two bond pairs around it, but no lone pair.

AB₃ type (Trigonal planar geometry)
In AB₃ type molecules, the central atom is surrounded by three electron pairs, which are arranged at maximum distance apart at an angle of 120° giving a trigonal geometry.
Example: BF₃ molecule has a trigonal planar shape with each F-B-F bond angle of 120° figure. The similar geometries are expected in the hydrides of group 3, i.e. AlH₃, GaH₃, BH₃, etc.

SnCl₂ In this molecule, the Sn atom has 4 electrons in its outermost shell. It makes two bonds with two Cl atoms and the remaining two electrons exist as a lone pair. One of the corners of the triangle is occupied by this lone pair, giving rise to a distorted trigonal electron pair geometry (in vapour phase). The actual shape of SnCl₂ is v-shaped and bond angle less than 120° due to the presence of a lone pair on Sn atom.

Ans. AB₄ Type (Tetrahedral geometry)
In this type of molecules the electron pair geometry is a regular tetrahedron. Each of the bond angles is of 109.5°.

(i) CH₄ (Methane)
In methane, four bonded electron pairs are directed from the center towards the corners of a regular tetrahedron, as following figure. On the same grounds, SiH₄, GeH₄, CCl₄ possess the similar shape.

(ii) NH₃ (Ammonia)
In the ammonia molecule (NH₃), there are three H atoms attached to the nitrogen atom having one lone pair. Due to a lone pair, the ideal angle 109.5° is reduced to 107° as in figure. Ammonia has triangular pyramidal shape. The substitution of hydrogen in NH₃ with electronegative atoms like F or Cl further reduces the bond angles.

(iii) NF₃ (Nitrogen trifluoride)
In NF₃, three highly electronegative F atoms are bonded to the N atom. The first reason being the strong polarity of N-F bond, due to which N atom pulls the lone pair closer to its nucleus. The lone pair exerts a stronger repulsion on bonding electrons. Moreover, the bond pairs (N-F bonds) are closer to F atoms than N atom. The increased distances in these bond pairs make their repulsions weaker and allow the bonds to come closer. The bond angles are further compressed to 102.5°.

(iv) H₂O (Water)
The water molecule (H₂O) is angular or v-shaped, although it has four electron pairs around the central atom. Two of the corners of the tetrahedon are occupied by two lone pairs and the remaining two by bond pairs. But, due to the greater repulsions of the lone pairs, the bond angle is reduced to 104.5° as in figure.

(v) H₃O⁺ (Hydronium ion) and NH₂⁻ (amide ion)
The [H₃O⁺] ion is formed by the capture of a proton by a water molecule and it has the tetrahedral trigonal pyramidal shape. On the other hand, amid ion, [NH₂]⁻, has a V–shaped structure as in figure.

Ans. AB₅ type molecules (trigonal bipyramidal geometry)
AB₅ type molecules, have trigonal bipyramidal geometry.

PCl₅
• This is evident from the geometry of the PCl₅ molecule given in figure.
• The angles within the plane are 120° and that between the axial atom and a planar atom is 90°.

Tri-iodide ion [I₃]⁻
• It is an example of AB₅ system.
• The central atom in this ion is the iodine atom.
• There are five electron pairs around the iodine atom including two bond pairs and three lone pairs.
• The lone pairs occupy the corners of the trigonal region and the iodine atoms occur vertically one being above and the second below the plane.
• The actual shape of the tri-iodide ion is linear and the iodine atoms are farthest apart at an angle of 180° as in figure.

AB₆ type molecules (octahedral geometry)
AB₆ type molecules have 6 electron pairs around the central atom and the octahedral geometry. All the ABA angles in this type of geometry are of 90°.
SF₆ is an AB₆ type molecule. The Lewis structure and the shape of SF₆ is given in figure.

XeF₄ is an AB₆ type molecule with four bond pairs and two lone pairs. the two lone pairs remain at farthest distance at angle of 180°. Excluding the lone pairs, the actual shape of XeF₄ is square planar with all F-P-F angles equal to 90° as in figure.

Ans. Applications of VSEPR Model in Drug Designing

Definition of drug Drugs are chemical substances that prevent, diagnose, or treat a disease.

• Drugs interact with human body at specific points, or with specific processes called targets or substrates, through their bioactive molecules called ligands.
• A receptor is mostly a protein that receives and responds to a ligand through binding.
• The examples of targets in the body are receptors, such as enzymes, proteins, nucleic acids, and cellular pathways.
• Each drug interacts with a specific target in the body to a different degree, this feature is called specificity of a drug or a ligand.

Molecular shape is an important feature that determines how a drug interacts with a biological target. Only the ligands with suitable shape can fit in the active sites of a biological target as a specific key fits in a specific lock. VSEPR model is successfully applied in determining the shapes of various biological systems, such as substrate recognition, ligand specificity or selectivity, and antibody recognition.

Aspirin
• It is an analgesic drug (painkiller) used for relief from pain as a primary medicine.
• It interacts with an enzyme COX (cyclooxygenase) by binding to its active site through the acetyl group (-OCCH₃).
• The shape of COX active site and that of acetyl group on Aspirin are compatible with each other as in figure. The binding is successful to block COX, which is a cause of the pain in the body.

Ans. Postulates of MOT
The molecules orbital approach explains the results of quantum mechanical calculations for covalent bonding.

The postulates of this theory:
(i) Overlapping of orbitals: During the formation of a molecule, the atomic orbitals of the combining atoms overlap to form new orbitals called 'molecular orbitals', which are characteristic of the whole molecule.
(ii) Polarization of orbitals: The atomic orbitals overlap with the lobe having suitable sign of the wave function of the orbital. For example, one lobe of a p orbital is given the + sign and the other is marked with '−'sign.
(iii) Types of molecular orbitals: Two atomic orbitals overlap to form two molecular orbitals. When same sign orbitals overlap, the bonding molecular orbital (σ or π) is formed that has lower energy than the parent atomic orbital, while with opposite signs, high energy anti-bonding molecular orbital (σ * or π*) is formed, that has higher energy than the parent atomic orbitals.
(iv) Bond order: 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

s-s Overlap:
H₂ molecule: In the formation of H₂ molecule, two s orbitals of H atoms combine to give two molecular orbitals. The bonding molecular orbital is symmetrical about the axis joining the nuclei of the bonded atoms (nuclear axis), while the anti-bonding molecular orbital has the electron density away from the nuclei of the overlapping atoms.

Ans. Molecular orbital diagram of H₂
• In the formation of H₂ molecule, two s orbitals of H atoms combine to give two molecular orbitals.
• The bonding molecular orbital is symmetrical about the axis joining the nuclei of the bonded atoms (nuclear axis), while the anti-bonding molecular orbital has the electron density away from the nuclei of the overlapping atoms.
• One of these molecular orbitals is the low-energy bonding molecular orbital (σ 1s) as there are only two electrons to be accommodated, these enter the lower energy σ 1s B.M.O. σ* 1s remains empty.
• The electrons in the bonding molecular orbital have opposite spin as shown in figure.

Bond Order = (No. of Electrons in B.M.O) - (No. of Electrons in A.B.M.O) = 2 - 0 = 1
2 2

2H(1s¹) → H₂(σ 1s²)

(ii) Molecular orbital diagram of He (Helium)
The electron configuration of He is 1s².
For a successful formation 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.
• Out of four electrons, two enter the bonding molecular orbital (σ 1s), and the remaining two occupy the antibonding (σ* 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 = 0
2

(iii) Molecular orbital diagram of Nitrogen (N₂)
Electronic configuration of N is N = 1s², 2s², 2p_x¹, 2p_y¹, 2p_z¹.
The molecular orbital diagram of N₂ based on this electron configuration is shown in the figure below.
The valence shell 2s on both N atoms give σ 2s, σ* 2s orbitals; whereas, 2p orbitals give six molecular orbitals give six molecular orbitals which are arranged in the increasing order of energy as:
σ 2s < σ* 2s < π2p_y = π2p_z < σ(2p_z) < π* 2p_y = π* 2p_z < σ* 2p_x

The bond order of N₂ can be calculated from its orbital diagram as:

Bond order of N₂ = 6 - 0 = 3
2

There are three covalent bonds between the nitrogen atoms in this molecule. One of these is a σ bond while the other two are π bonds.
Nitrogen molecule has a triple bond between its atoms.

(iv) continues on next page...

Ans. (i) Lewis structure of F₂

:F· - F:
xx
xx

(ii) Structure of F₂ by VBT
Formation of F₂ according to VBT p-orbitals of two fluorine atoms overlap on the nuclear is the electron density is symmetrical' around the nuclei of two fluorine because both have same electronegativity value.

(iii) Formation of F₂ according to MOT
⁹F = 1s², 2s², 2p⁵

The molecular orbital diagram of F₂ based on this electron configuration is shown in the figure.

Bond order of F₂ = 6 - 4 = 2 = 1
2 2
Bond order of F₂⁺ = 1