Class 11 Chemical Bonding Notes
CHEMICAL BOND
It is the force of attraction between two atoms which hold them together in a compound or molecule.
Why Chemical Bond Formation Take Place?
Nature loves stability and bond formation is associated with stability. An element has a tendency to occupy inert electronic configuration (ns2 np6 or ns2) which is considered as very stable. Noble gas electronic configuration can be achieved by
(1) Transference of electrons
(2) Mutual sharing of electrons
(3) Donation of lone pair of electrons
Types of Bond
1. lonic bond or electrovalent bond
2 Covalent bond
3 Co-ordinate bond
4 Metallic bond
5. Hydrogen bond
6 van der Waals bond
In order to explain the formation of a chemical bond in terms of electrons, a number of attempts were made.
Lewis postulated that atoms achieve stable octet when they are linked by a chemical bond .
LEWIS DOT STRUCTURES
In the formation of a molecule only the outer shell electrons take part in chemical bond combination and they are known as valence electrons.
In Lewis symbols, an element is shown with symbol and valence electrons are shown around with a dot.
Every atom has a tendency to attain Noble gas electronic configuration or to have 8 valence electrons. This is known as law of octet rule or if it has two valence electrons then this is known as law of duplet. According to Lewis, only those compounds will be stable which follow octet rule .
Limitation of Lewis Octet Rule
(1) Incomplete octet of the central atom There are some compounds in which central atom does not have complete octet but still stable.
(3) Xe which is a noble gas but still form some compounds .
(4) Lewis octet rule gave no idea about molecular shape of the compound
Lewis dot representations provide a picture of bonding in molecules and ions in terms of shared par of electrons and the octet rule. Although it does not explain the complete picture of bonding but it helps in understanding the formation of molecule Rules which are adopted to write Lewis dot structure for molecule are as
1 ) Calculate the total number of valence electrons available for bonding in given molecule
For anion à Add electrons equal to charge in valence electrons of that atom
For cation à Remove electrons equal to charge in valence electrons of that atom
For NH4+ → Number of valence electrons = 5 + 4 – 1 = 8
2 ) Least electronegative polyvalent atom occupies central position.
3) Show shared pair of electrons between two atoms and unshared electrons on one atom
Formal Charge
The molecule is neutral and its constituent atom do not carry any charge similarly in case of polyatomic ions the net charge is possessed by the ion as a whole and not by a particular atom. Yet for some purposes it is very useful to assign formal charges to the individual atoms constituting molecules or even polyatomic ions
Formal charge on an atom is the difference between the number of valence electrons is an isolated atom (ie free atom) and the number of electrons assigned to that atom in a Lewis structure It is expressed as :
Formal charge = V – L - 1/2S
The counting is based on the assumption, that the atom in the molecule possess one electron of each shared pair and both the electrons of a lone pair
IONIC BOND
An ionic bond is formed by complete transference of one or more electrons from the valence shell of one atom to the valence shell of another atom. In this way both the atoms acquire stable electronic configurations of noble gases The atom which loses electron becomes a positive ion and the atom which gains electron becomes negative ion
The electrostatic force of attraction between the oppositely charged ions result in the formation of an ionic bond
Na à Na+ + e-
Factors Influencing the Formation of Ionic Bond
The following factors influence the formation of ionic bond
1. Ionization Energy: It is defined as the energy required to remove the most loosely bound electron from an isolated gaseous atom of an element. The lesser is the ionization energy greater is the ease of formation of a cation
2. Electron Affinity: It is defined as the amount of energy released when an electron is added to an isolated gaseous atom of an element
The formation of a non-metal anion occurs with the addition of one or more electrons to the non-metal atom. The higher the energy released during this process, the easier will be the formation of an anion .
3) Lattice Energy : It is defined as the amount of energy released when gaseous cations and anions are brought from infinity to their equilibrium sites in the crystal lattice, to form one mole of an ionic compound.
As lattice energy is an ionic force. so according to Coulomb's law of attraction
F = -Kq1q2 / r2
Characteristics of lonic Compounds
1) They are hard, brittle and crystalline
2) They have high melting and boiling points.
3)They are polar in nature.
4) The linkage between oppositely charged ions is non rigid and non directional
5) They are soluble in polar solvents such as water and insoluble in non polar solvents such as CCI, Benzene, ether etc
(a) Water has a high dielectric constant (80), therefore, ionic solids when kept in water, the forces of attraction among ions are reduced to 1/80 and ions are free to pass in solution
(b) Water is polar, therefore, dipole ion attraction leads to separation of ions and makes them pass into solution
(c) lons have a tendency to get hydrated and the energy released is known as hydration energy. If hydration energy is greater than lattice energy then ions will fall apart in solution, Le lonic compound becomes soluble.
6) They are good conductors of electricity in fused state and in solution due to mobility of the ions. They are bad conductors of electricity in solid state because ions are unable to move.
7) Since they form ions in solution hence they exhibit ionic reactions which are quite faster and instantaneous
COVALENT BOND
The gain or loss of electrons cannot take place between similar atoms. In such cases the band formed by mutual sharing of electrons. A force which binds atoms of same or different elements by mutual sharing of electrons is called a covalent bond
Each chlorine atom, Ci (2. 8, 7 or [Ne]3s 3phy has seven electrons in its valence-shell and needs one more electron to acquire octet or to attain the electronic configuration of Argon. Both the chlorine atoms contribute one electron each to share two electrons (shared pair) The bond can be represented by putting a line (-) between the atoms or with electron dot symbols also.
Molecules containing identical or different atom can also be represented with the Lewis dot structures there are some important condition
1) Bond should be formed by the sharing of an electron pair between the atoms
2) 2) There should be a contribution of at least one electron by each combining atom to the shared pair.
3) 3)The sharing of electrons should be in such a way that the combining atoms attain the outer shell noble gas configurations .
VALENCE BOND THEORY (VBT)
This approach was developed by Heitler and London in 1927 and further improved by Pauling and others Understanding VBT involves the knowledge of atomic orbitals, electronic configurations of elements, the overlap criteria of atomic orbitals and the principles of variation and superposition.
It is well known fact that all the mechanical systems in this universe tend to lower their potential energy so that they can attain the greater stability Same is the situation involved with the formation of bonds between the atoms, bonding occurs with decrease in energy. In order to understand this concepts let us study the formation of H, molecule on the basis of electrostatic interactions which lead to decrease of energy Attractive forces tend to bring the two atoms close to each other whereas repulsive forces tend to push them apart.
The magnitude of new attractive forces is more than the new repulsive forces. As a result, two atoms approach each other and potential energy decreases Ultimately, a state is reached where attractive forces, just balance the repulsive forces. This state is reached when the atoms are at critical distance r At a distance greater than r, the attractive forces are dominant whereas at a distance smaller than r, the repulsive forces dominate over the attractive forces which leads to increase of potential energy Maximum lowering of energy takes place at critical distance. The lowering of PE as a function of inter nuclear distance has been shown in figure below
It is clear from the curve that the energy of two H atoms when they are held at a distance r, is smaller than the energy of individual H atoms. Therefore, the H atoms constitute a stable grouping called hydrogen molecule. The critical distancer, corresponding to minimum energy is called bond length Experiments have shown that the value of r, is 74 pm
Orbital overlap concept
If we refer to the minimum energy state in the formation of hydrogen molecule the two H-atoms are enough near so as to allow their atomic orbitals to undergo partial interpenetration. This partial interpenetration of atomic orbitals is called overlapping of atomic orbitals
The overlap between the atomic orbitals can be positive, negative or zero depending upon the characteristics of the orbitals participating to overlap
Positive overlap involves the overlap of the lobes of same signs. It leads to attractive interactions.
Negative overlap imvolves the overlap of the lobes of opposite signs it leads to repulsive interactions Zero overlap implies inability of any kind of interactions. The various arrangements of s and p orbitals resulting in positive, negative and zero overlap are depicted in figure below :
As we have already discussed that the formation of a covalent bond involves the overlapping of half filled atomic orbitals. The covalent bonds can be classified into two different categories depending upon the type of overlapping These are :
1 ) Sigma bond and
2 ) pi bond
1. Sigma (a) bond : This type of covalent bond is formed by the axial overlapping of half-filled atomic orbitals. The atomic orbitals overlap along the internuclear axis and involve end to end or head on overlap . There can be three type of aut overlap among s and p-orbitals as discussed below
(1) s-s overlap : In this case there is overlap of two half filled a-orbitals along the internuclear as shown below
(ii) s-p overlapping : It involves the overlapping of hall filled 5-orbitals of one atom with the half filled p-orbitals of the other atom. The bond thus formed is called s-p sigma bond .
For example, the formation of HF molecule involves the overlapping of 13-orbital of H-atom with the half filled 2p, orbital of fluorine atom.
(iii) p-p overlapping : it involves the co-axial overlapping between half filled p-orbitals of one atom with half filled p-orbitals of the other atom. The bond as formed is called p-p sigma bond
HYBRIDISATION
In order to explain the characteristic geometrical shapes of polyatomic molecules like CH, NH, and H.O elc a concept called hybridisation was introduced by Linus Pauling According to him the atomic orbitals combine to form new set of equivalent orbitals known as hybrid orbitals or hybridised orbitals and the phenomenon is referred to as hybridisation
Hybridisation is the process of intermoxing 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.
Salient Features of Hybridisation
Hybridisation is a theoretical concept which has been introduced to explain some structural properties such as shapes of molecules or equivalency of bonds, etc., which cannot be explained by simple theories of valency Some salient-features of hybridisation are
1 The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised
2 The hybridised orbitals are always equivalent in energy and shape.
3 The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals
4. These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement. Therefore the type of hybridisation indicates the geometry of the molecules .
Important conditions for hybridisation
(i) The orbitals present in the valence shell of the atom are hybridised
(ii) The orbitals taking part in hybridisation must have only a small difference of energies
(iii) Promotion of electron is not essential condition prior to hybridisation
(iv) It is not necessary that only half filled orbitals participate in hybridisation In some cases, even filled orbitals of valence shell or vacant orbital can take part in hybridisation
Types of hybridisation
There are many different types of hybridisation depending upon the type of orbitals involved in muxing such as sp3 sp2 sp sp3d2 of etc. Let us discuss various types of hybridisation along with some examples
(1) sp-hybridisation
In this hybridisation one s and one p orbitals hybridise (or intermix) to produce two equivalent hybird orbitals, known as sp hybrid orbitals. The suitable orbitals for sp hybridisation are s and p, if the hybrid orbitals are to lie along the z-axis. The two sp-hybrid orbitals are oriented in a straight line making an angle of 180" and therefore the molecule possesses linear geometry Each of hybrid orbitals has 50% s-character and 50% p-character This type of hybridisation is also known as diagonal hybridisation
II) sp2 hybridisation
Sp2 hybrid orbitals are larger in size than sp-hybrid orbital but slightly smaller than that of sp3 hybrid orbitals. Each sp hybrid orbitals has 1/3 (or 33.33%) s-character and 2/3 (or 667%) p-character
(II) sp3 hybridisation
In this hybridisation ones and three p-orbitals intermix to form sp hybrid orbitals of equivalent energy and identical shape. These four sp³ hybrid orbitals are directed towards the four comers of a tetrahedron separated by an angle of 109" 28 sp hybrid orbitals have 1/4 (or 25%) s-character and 3/4 (or 75%) p-character .
PREDICTION OF STATE OF HYBRIDISATION
Various methods are available to determine state of hybridisation However most appropriate one is electron pair (EP) method, that is described as under
Count electron pairs as
EP = BP + LP
Valence shell electron pair repulsion theory
(Developed by Gillespie and Nyholm)
1 ) According to this theory, shape of the compound depends on the number and types of electron pairs around the central atom in compound.
(a) Bond pair Pair of electrons under the influence of two atoms
(b) Lone pair Pair of electrons localised on one atom
2) These electron pairs repel each other and stay as far as possible. Sequence of repulsion is
(a) Lone pair-lone pair> Lone pair-bond pair> Bond pair-bond pair
3) Double bond gives more repulsion than single bond.
COVALENT CHARACTER IN IONIC BONDS
When oppositely charged lons approach each other, there is not only the attraction between the positively charged cation and the negatively charged anion but also simultaneous repulsions between their nuclei. Thus there is distortion, or deformation or polarization of anions. The electronic charge of anion does not remain spherical but gets distorted
Fajan's Rule
Covalent character in lonic compound
There is no compound which is 100% lonic Le every lonic compound has some covalent character
Induction of covalent character ie. polarisation of ionic compound is a result of distortion of anion by a cation which is known as polarisation of anion by a cation. More is the polarisation (distortion of anion), more will be the percentage of covalent character Covalent character is relatively measured with the help of Fajan's rule:
1) Smaller cation has more polarising power, because of high ve charge density and effective nuclear charge. Because of which smaller caton has a stronger hold on the e's of nearby anion, that leads to distorsion of anion's e cloud and introduces covalent character.
2) Larger anions have high polarisability because, in large anion, nucleus of anion will not have as much strong hold on outer electrons as that of smaller anion, ie larger anions are more prone to be distorted, hence, more covalent character in the compound with larger size anions
3) Larger the magnitude of charge on cation and anion more is the polarising power of cation and more. is the tendency of anion to get polarised, hence larger is the covalent character
4) . Cation with 18 electrons in outermost shell Le with pseudo noble gas electronic configuration will have more polarising power than that of a cation with 8 electrons in outer most shell (noble gas type electronic configuration)
Some important applications of covalent character in an ionic compound
i) Melting point decreases with increase in covalent character
(ii) Solubility decreases in polar solvent (H,O) and increases in non polar solvents as covalent character increases .
RESONANCE
When a compound has same molecular formula but different structures differ with respect to electrons only. These structures are known as resonating structures or canonical structures None of these structures can explain all the properties of that compound. This phenomenon is known as resonance
Bond order in a compound showing resonating structure = total number of bonds / total number of resonating structures
Resonance Hybrid
It is the hypothetical mixture of all contributing or resonating structure
Conditions for Writing a Resonating Structure
1 ) Maximum valency of an atom should not be violated
2) Sum of number of t-bonds and lone pair of electrons in every structure must be same
MOLECULAR ORBITAL THEORY (MOT)
The quantum wave mechanical treatment of a covalent bond as based on the valence bond theory is discussed earlier The second approach to the problem is through another theory, known as molecular orbital theory The two theories reflect distinctly different conceptual approaches to the basic structural model of a molecule For instance, according to VBT the two atoms in a diatomic molecule come close together while their completely filled orbitals remains intact. Only half filled orbitals of one atom overlaps with the half filled atomic orbitals of the other atom, resulting in the formation of a covalent bond
According to MOT a molecule is considered to be quite different from the constituent atomsAll the electrons belonging to the atoms constituting a molecule are considered to be moving along the entire molecule under the influence of all the nuclei Thus, a molecule is supposed to have orbitals of varying energy levels, in same way as an atom. These are called molecular orbitals In other words while atomic orbitals are monocentricmolecular orbitals are polycentric
The main ideas of this theory can be summed up as follows
1) Just like VBT it also makes use of atomic orbitals, but atomic orbitals of one atom overlap with orbitals atomic orbitals of other atom to form entirely new orbitals called molecular orbitals where the atomic orbitals loose their individual identity .
2) Molecular orbitals just like atomic orbitalsare energy states of the molecule in which electrons of the molecule are filled.
3) A molecular orbital gives the electron probability distribution around a group of atoms or molecules.