Chapter 4 : Chemical bonding and molecular structure
According to Kossel-lewis the chemical bond is:
Topics covered in this snack-sized chapter:
- The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the electrovalent bond.
According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.
A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms.
- The electrovalence is thus equal to the number of unit charge(s) on the ion.
The formation of the Cl2
molecule can be understood in terms of the sharing of a pair of electrons between the two chlorine atoms, each chlorine atom contributing one electron to the shared pair.
When two atoms share one electron pair they are said to be joined by a single covalent bond.
If two atoms share two pairs of electrons, the covalent bond between them is called a double bond.
Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule.
Lewis structures show each atom and its position in the structure of the molecule using its chemical symbol.
4.1.3 Lewis Representation of Simple
Molecules (the Lewis structures)
A formal charge (FC) is the charge assigned to an atom in a molecule, assuming that electrons in a chemical bond are shared equally between atoms, regardless of relative electronegativity.
4.1.5 Limitations of the Octet Rule
The incomplete octet of the central atom
In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are LiCl, BeH2
In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO2
, the octet rule is not satisfied for all the atoms.
The expanded octet:
Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet.
Obviously the octet rule does not apply in such cases. Some of the examples of such compounds are: PF5
An electrovalent bond is a chemical bond in which one atom loses an electron to form a positive ion and the other atom gains an electron to form a negative ion.
The formation of a positive ion involves ionization, i.e., removal of electron(s) from the neutral atom and that of the negative ion involves the addition of electron(s) to the neutral atom.
(g) + e–
X(g) + e–
Electron gain enthalpy
(g) + X–
(g) → MX(s)
Obviously ionic bonds will be formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy.
Ionic compounds in the crystalline state consist of orderly three-dimensional arrangements of cations and anions.
The crystal structure of sodium chloride, NaCl (rock salt), for example is shown below:
In ionic solids, the sum of the electron gain enthalpy and the ionization enthalpy may be positive but still the crystal structure gets stabilized due to the energy released in the formation of the crystal lattice.
The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.
For example, the lattice enthalpy of NaCl is 788 kJ mol–1
. This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+
(g) and one mole of Cl–
(g) to an infinite distance.
Some bond parameters are:
Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule.
Bond Angle is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. It is expressed in degree.
It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.
The unit of bond enthalpy is kJ mol-1
. For example, the H ñ H bond enthalpy in hydrogen molecule is 435.8 kJ mol-1
Bond Order is given by the number of bonds between the two atoms in a molecule. The bond order, for example in H2
(with a single shared electron pair), in O2
(with two shared electron pairs) and in N2
(with three shared electron pairs) is 1, 2, 3 respectively.
With increase in bond order, bond enthalpy increases and bond length decreases.
According to the concept of resonance, whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.
Thus for O3
, the two structures shown above constitute the canonical structures or resonance structures and their hybrid i.e., the III structure represents the structure of O3
more accurately. This is also called resonance hybrid. Resonance is represented by a double headed arrow.
4.3.5 Resonance Structures
In general, it may be stated that
Resonance averages the bond characteristics as a whole.
Thus the energy of the O3
resonance hybrid is lower than either of the two canonical forms I and II.
The existence of a hundred percent ionic or covalent bond represents an ideal situation.
In reality no bond or a compound is either completely covalent or ionic.
In case of HF, the shared electron pair between the two atoms gets displaced more towards fluorine since the electronegativity of is far greater than that of hydrogen. The resultant covalent bond is a polar covalent bond.
As a result of polarization, the molecule possesses the dipole moment.
The main postulates of VSEPR theory are as follows:
4.4 The Valence Shell Electron Pair Repulsion (VSEPR) Theory arrow_upward
- The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom.
- Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
- These pairs of electrons tend to occupy such positions in space that minimize repulsion and thus maximize distance between them.
- The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at maximum distance from one another.
- A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair.
- Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure.
- The repulsive interaction of electron pairs decrease in the order:
Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp)
- The lone pairs are localized on the central atom, each bonded pair is shared between two atoms.
Consider two hydrogen atoms A and B approaching each other having nuclei NA
and electrons present in them are represented by eA
. When the two atoms are at large distance from each other, there is no interaction between them. As these two atoms approach each other, new attractive and repulsive forces begin to operate.
Attractive forces arise between:
- The VSEPR Theory is able to predict geometry of a large number of molecules, especially the compounds of p-block elements accurately.
- Nucleus of one atom and its own electron that is NA
Similarly repulsive forces arise between
- Nucleus of one atom and electron of other atom i.e., NA
, NB – eA
- Electrons of two atoms like eA
Attractive forces tend to bring the two atoms close to each other whereas repulsive forces tend to push them apart.
Experimentally it has been found that the magnitude of new attractive force is more than the new repulsive forces. As a result, two atoms approach each other and potential energy decreases.
Ultimately a stage is reached where the net force of attraction balances the force of repulsion and system acquires minimum energy. At this stage two hydrogen atoms are said to be bonded together to form a stable molecule having the bond length of 74 pm.
Since the energy gets released when the bond is formed between two hydrogen atoms, the hydrogen molecule is more stable than that of isolated hydrogen atoms. The energy so released is called as bond enthalpy, which is corresponding to minimum in the curve depicted in the figure below. Conversely, 435.8 kJ of energy is required to dissociate one mole of H2
- Nuclei of two atoms NA
(g) + 435.8 kJ mol-1
H(g) + H(g)
In the formation of hydrogen molecule, there is a minimum energy state when two hydrogen atoms are so near that their atomic orbitals undergo partial interpenetration. This partial merging of atomic orbitals is called overlapping of atomic orbitals which results in the pairing of electrons.
Greater the overlap the stronger is the bond formed between two atoms.
According to orbital overlap concept, the formation of a covalent bond between two atoms results by pairing of electrons present in the valence shell having opposite spins.
4.5.1 Orbital Overlap Concept:
When orbitals of two atoms come close to form bond, their overlap may be positive, negative or zero depending upon the sign (phase) and direction of orientation of amplitude of orbital wave function in space.
Various overlaps of s and p orbitals are depicted in the figure given below:
4.5.3 Overlapping of Atomic Orbitals:
Covalent Bonds are classified into two types:
4.5.4 Types of Overlapping and Nature of Covalent Bonds arrow_upward
- Sigma bond, and
- pi bond
This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the inter-nuclear axis. This is called as head on overlap or axial overlap.
In this case, there is overlap of two half filled s-orbitals along the inter-nuclear axis as shown below:
This type of overlap occurs between half-filled s-orbitals of one atom and half-filled p-orbitals of another atom.
This type of overlap takes place between half filled p-orbitals of the two approaching atoms.
In the formation of bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the inter-nuclear axis. The orbitals formed due to sidewise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms.
The strength of a bond depends upon the extent of overlapping. In case of sigma bond, the overlapping of orbitals takes place to a larger extent. Hence, it is stronger as compared to the pi bond where the extent of overlapping occurs to a smaller extent. Further, it is important to note that in the formation of multiple bonds between two atoms of a molecule, pi bond(s) is formed in addition to a sigma bond.
Hybridization is defined as the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape.
For example when one 2s and three 2p-orbitals of carbon hybridise, there is the formation of four new sp3
The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridized.
The hybridized orbitals are always equivalent in energy and shape.
The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.
These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement.
Salient features of hybridization:
This type of hybridisation involves the mixing of one s and one p orbital resulting in the formation of two equivalent sp hybrid orbitals.
Such a molecule in which the central atom is sp-hybridised and linked directly to two other central atoms possesses linear geometry. This type of hybridisation is also known as diagonal hybridisation.
Example of molecule having sp hybridization BeCl2
The ground state electronic configuration of Be is 1s2
. In the exited state one of the 2s-electrons is promoted to vacant 2p orbital to account for its bivalency.
One 2s and one 2p-orbital gets hybridised to form two sp hybridised orbitals. These two sp hybrid orbitals are oriented in opposite direction forming an angle of 180°. Each of the sp hybridised orbital overlaps with the 2p-orbital of chlorine axially and form two Be- Cl sigma bonds.
In this hybridization there is involvement of one s and two p-orbitals in order to form three equivalent sp2
hybridised orbitals. For example, in BCl3
molecule, the ground state electronic configuration of central boron atom is 1s2
In the excited state, one of the 2s electrons is promoted to vacant 2p orbital as a result boron has three unpaired electrons.
These three orbitals (one 2s and two 2p) hybridise to form three sp2
hybrid orbitals. The three hybrid orbitals so formed are oriented in a trigonal planar arrangement and overlap with 2p orbitals of chlorine to form three B-Cl bonds.
This type of hybridization can be explained by taking the example of CH4
molecule in which there is mixing of one s-orbital and three p-orbitals of the valence shell to form four sp3
hybrid orbital of equivalent energies and shape.
The four sp3
hybrid orbitals so formed are directed towards the four corners of the tetrahedron. The angle between sp3
hybrid orbital is 109.5.
Hybridisation in C2
molecule: In ethane molecule both the carbon atoms assume sp3 hybrid state
Hybridisation in C2
: In the formation of ethene molecule, one of the sp2
hybrid orbitals of carbon atom overlaps axially with sp2
hybridised orbital of another carbon atom.
sp Hybridisation in C2
: In the formation of ethyne molecule, both the carbon atoms undergo sp-hybridisation having two unhybridised orbital i.e., 2py
4.6.2 Other Examples of sp3
The hybridisation involving either 3s, 3p and 3d or 3d, 4s and 4p is possible.
The important hybridisation schemes involving s, p and d orbitals are summarized below:
4.6.3 Hybridisation of Elements involving d Orbitals
Formation of PCl5
The ground state and the excited state outer electronic configurations of phosphorus (Z=15).
Now the five orbitals (i.e., one s, three p and one d orbitals) are available for hybridisation to yield a set of five sp3d hybrid orbitals which are directed towards the five corners of a trigonal bipyramidal as depicted in the figure given below:
Formation of SF6
the central sulphur atom has the ground state outer electronic configuration 3s2
In the excited state the available six orbitals i.e., one s, three p and two d are singly occupied by electrons.
These orbitals hybridise to form six new sp3
hybrid orbitals, which are projected towards the six corners of a regular octahedron in SF6
These six sp3d2
hybrid orbitals overlap with singly occupied orbitals of fluorine atoms to form six S-F sigma bonds. Thus SF6
molecule has a regular octahedral geometry as shown in the figure below.
The salient features of this theory are:
- The electrons in a molecule are present in the various molecular orbitals as the electrons of atoms are present in the various atomic orbitals.
- The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals.
- While an electron in an atomic orbital is influenced by one nucleus, in a molecular orbital it is influenced by two or more nuclei depending upon the number of atoms in the molecule. Thus, an atomic orbital is mono-centric while a molecular orbital is polycentric.
- The molecular orbitals like atomic orbitals are filled in accordance with the aufbau principle obeying the Pauli’s exclusion principle and the Hund’s rule.
4.7.1 Formation of Molecular Orbitals Linear Combination of Atomic Orbitals (LCAO)
Consider the hydrogen molecule consisting of two atoms A and B. Each hydrogen atom in the ground state has one electron in 1s orbital. The atomic orbitals of these atoms may be represented by the wave functions and .
Mathematically, the formation of molecular orbitals may be described by the linear combination of atomic orbitals that can take place by addition and by subtraction of wave functions of individual atomic orbitals as shown below:
Therefore, the two molecular orbitals and are formed as:
The molecular orbital formed by the addition of atomic orbitals is called the bonding molecular orbital while the molecular orbital formed by the subtraction of atomic orbital is called antibonding molecular orbital.
The linear combination of atomic orbitals to form molecular orbitals takes place only if the following conditions are satisfied:
4.7.2 Conditions for the Combination of Atomic Orbitals
- The combining atomic orbitals must have the same or nearly the same energy.
- The combining atomic orbitals must have the same symmetry about the molecular axis.
- The combining atomic orbitals must overlap to the maximum extent.
4.7.3 Types of Molecular Orbitals
Molecular orbitals of diatomic molecules are designated as (sigma), (pi), (delta), etc.
The sigma molecular orbitals are symmetrical around the bond-axis while pi molecular orbitals are not symmetrical.
For example, the linear combination of 1s orbitals centered on two nuclei produces two molecular orbitals which are symmetrical around the bond-axis. Such molecular orbitals are of the σ type and are designated as and . If inter nuclear axis is taken to be in the z-direction, it can be seen that a linear combination of 2pz
- orbitals of two atoms also produces two sigma molecular orbitals designated as 2pz
Molecular orbitals obtained from 2px
orbitals are not symmetrical around the bond axis because of the presence of positive lobes above and negative lobes below the molecular plane. Such molecular orbitals, are labelled as and .
atomic orbitals on two atoms form two molecular orbitals designated as σ1s
In the same manner, the 2s
The increasing order of energies of various molecular orbitals is:
4.7.4 Energy Level Diagram for Molecular Orbitals
The distribution of electrons among various molecular orbitals is called the electronic configuration of the molecule.
4.7.5 Electronic Configuration and Molecular Behavior
Stability of Molecules:
is the number of electrons occupying bonding orbitals and Na
the number occupying the anti-bonding orbitals, then
- The molecule is stable if Nb
is greater than Na
- The molecule is unstable if Nb
is less than Na
Bond order (b.o.) is defined as one half the differences between the number of electrons present in the bonding and the anti-bonding orbitals i.e.
Bond order (b.o.) = (1/2)(Nb
Nature of the bond:
Integral bond order values of 1, 2 or 3 correspond to single, double or triple bonds respectively.
The bond order between two atoms in a molecule may be taken as an approximate measure of the bond length. The bond length decreases as bond order increases.
If all the molecular orbitals in a molecule are doubly occupied, the substance is diamagnetic. However if one or more molecular orbitals are singly occupied it is paramagnetic.
Hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule.
Hydrogen is bonded to strongly electronegative element because the electron pair shared between the two atoms moves far away from hydrogen atom
As a result the hydrogen atom becomes highly electropositive with respect to the other atom
4.9.1 Cause of Formation of Hydrogen Bond
There are two types of H-bonds
Intermolecular hydrogen bond:
It is formed between two different molecules of the same or different compounds. For example, H-bond in case of HF molecule, alcohol or water molecules, etc.
Intramolecular hydrogen bond:
It is formed when hydrogen atom is in between the two highly electronegative (F, O, N) atoms present within the same molecule. For example, in o
-nitrophenol the hydrogen is in between the two oxygen atoms.