Plus One Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure is part of Plus One Chemistry Notes. Here we have given Kerala Plus One Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure.
|Text Book||NCERT Based|
|Chapter Name||Chemical Bonding and Molecular Structure|
|Category||Plus One Kerala|
Kerala Plus One Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure
A group of atoms is found to exist together as one species having characteristic properties. Such a group of atoms is called a molecule. There must be some force which holds these constituent atoms together in the molecules.
The attractive force that binds the atoms together in a molecule is called a chemical bond,
eg., bond in molecule.
Kossel-Lewis Approach to Chemical Bonding
Lewis introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols.
The valence electrons are shown by electron dots or crosses around the symbol.
It indicates the positive or negative valency of the atom.
Covalent Bond, Langmuir in 1919 refined the Lewis postulations by abandoning the idea of the stationary cubical arrangement of the octet, and by introducing the term covalent bond.
By Lewis-Langmuir theory the information of chlorine molecule is as follows:
Formal charge (F.C) on an atom in a Lewis structure = total number of valence electrons in the free atom – total number of non bonding (lone pair) electrons – (1/2) total number of bonding (shared) electrons.
Electrovalent or Ionic bond
An ionic bond is formed by the complete transference of one or more valence electrons of one atom to the valence shell of the other atom.
The electrostatic attraction between the oppositely charged ions which always lead to a decrease in the potential energy of the system is known as ionic bond.
The 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 shell. This is known as octet rule.
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.
In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, N02, the octet rule is not satisfied for all the atoms.
The Expanded Octet
In a number of compounds of these elements there are more than eight valence electrons around the central atom.
Single bond and Multiple bond (Double bond, Triple bond)
When two atoms share one electron pair, the covalent bond between them is called a single bond,
eg., H-H bond.
When two atoms share two electron pair or three electron pair, the covalent bond between them is called multiple bond (double bond or triple bond),
eg., 0=0 bond, N≡ N bond.
Electron Gain Enthalpy (Δeg H)
It is the enthalpy change, when a gas phase atom in its ground state gains an electron, Lattice Enthalpy. It is defined as the amount of energy released when one mole of ionic solid is formed by the close packing of gaseous ions.
It is defined as the average distance between the centres of nuclei of the two bonded atoms in a molecule.
- Bond length increases with increase in the size of the atom.
- Bond length decreases with the multiplicity of bonds.
It is the half of the distance between two similar atoms joined by a covalent bond in the same molecule.
Van der Waal’s Radius.
It is half of the distance between two similar atoms in separate molecules in a solid.
It is defined as the average angle between the lines representing the orbitals containing the bonding electrons.
Bond Enthalpy or Bond Energy.
The average bond dissociation enthalpy required to dissociate each bond in a molecule is called bond enthalpy.
Bond dissociation enthalpy refers to the energy required to break one mole of a particular kind of bond in a molecule.
Factors affecting bond energy
- Atomic size α bond length
- Triple bond ( = ) > Double (=) > Single bond (-)
- Number of lone pairs of electrons
Bond Order (B.O).
It gives the number of covalent bonds between the two atoms in a molecule.
When a molecule cannot be represented by a single structure but its characteristic properties can be described by two or more than two structure, then the actual molecule is said to be a resonance hybrid of these structures.
0-0 = 148 pm, 0=0 = 121 pm. But experimentally all bonds in 03 have the same bond length 128 pm.
Predicting the Hybrid State of Atom in Different Species
- Write the valence electrons of central atom (V)
- Add number of surrounding atoms except Oxygen (SA).
- If there is positive charge on species, subtract the charge, and if there is negative charge, then add the charge (E).
- Divide the sum by 2 to get the value of X.
Mathematically, X = -1/2 (F + SA + E)
|Value of X||2||3||4||5||6||7|
|Hybrid state of central atom||sp||sp2||sp3||sp3d||sp3d2||sp3d3|
Note: The preceding formula is not applicable to predict hybrid state of metal in complexes and species having multicentral atoms.
Polarity Of Bonds
The pair of bonding electrons is not shared equally between two atoms. Such a bond is called polar or heteropolar covalent bond.
Dipole moment (μ)
It is defined as the product of the magnitude of the charge and the distance between the centres of positive and negative charge, dipole moment (μ) = charge (Q) x distance of separation (r)
VSEPR (Valence Shell Electron Pair Repulsion) Theory
Main postulates of VSEPR theory:
- The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom.
- Electron pair try to stay as far apart as possible to acquire a state of minimum energy or maximum stability.
- The presence of lone pairs in addition to bond pairs will result in distortions in the shape of the molecule.
- Repulsive force decrease sharply with increase in bond angle between electron pairs.
Valence Bond Theory: Formation of Hydrogen Molecule
Consider two hydrogen A and B approaching each other having nuclei NA and NB and electron present in them are represented by eA and eB . When the 2 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.
The atoms are held at critical distance γ0 when attractive forces balance the repulsive forces. When internuclear distance is greater than γ0, attractive force is dominant. When internuclear distance is smaller than γ0, the repulsive force is dominant.
Orbital Overlap Concept
According to this concept, a covalent bond is formed between two atoms when a half filled valence orbital of one atom overelaps with a half filled valence orbital of another. It is essential that the two electrons which are involved in bond formation should have their spins in opposite directions and get paired.
Overlapping of Atomic Orbitals
The overlapping of atomic orbitals along orbital axis or perpendicular to orbital axis may be positive overlap or negative overlap.
sigma (σ) bond (Head on overlap or axial overlap). When two half filled atomic orbitals overlap along the inter-nuclear axis (the line passing through the centres of nuclei of the two atoms) the bond formed is known as sigma bond. Sigma bond is strong, single bond and it has maximum overlapping.
- s-s overlapping
The mutual overlap of half filled s orbitals of two atoms always leads to the formation of a bond known as s-s bond.
- s-p overlapping
This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom.
- p-p overlapping
This type of overlap take place between half filled two p-orbitals.
eg., F2 molecule
Pi (π) bond.
When half filled p-orbitals involved in bond formation are oriented in a direction perpendicular to the internuclear axis, the two orbits overlap sidewise to form n bond, n bond is weak and containing double or triple bond.
It is defined as the phenomenon of intermixing of pure atomic orbitals of slightly different energies and shapes to form new hybrid orbitals of equivalent energies and identical shapes.
Types of Hybridisation: sp, sp2 and sp3 hybridisation
sp hybridisation or diagonal hybridisation.
In this type of hybridisation, only 2s and one of the 2p orbitals hybridise forming two equivalent orbitals. Each sp hybride orbitals has 50% s-character and 50% p character. Bond angle is 180° linear molecule,
eg., BeCI2, CH ≡ CH or acetylene.
Electronic configuration of Be (4) is 1s2 2s2 in ground state. It is Be(4) 1s2 2s1 2px1 in excited state
(since Be is divalent).
This hybridisation involves the mixing of one s and two p-orbitals in order to form three equivalent sp2 hybridised orbit.
Electronic configuration of B (5) is 1s2 2s2 2px1 in ground state. It 1133 electronic configuration 1s2 2s1 2px1 2py1 in excited state (since B is trivalent only in excited state).
In this hubridisation, there is involvent of one s-orbital an three p-orbitals of the valence shell to form four sp3 hybrid orbitals. Each sp3 hybrid orbitals has 25% s character and 75% p character.
Electronic configuration of C (6) is 1s2 2s2 2px1 2py1 in ground state. It has electronic configuration 1s2 2s1 2px1 2py1 2pz1 in excited state (since C is tetravalent only in excited state).
Molecular Orbital Theory
Molecular orbital (M.O) theory was developed by F. Hund and R. S. Mulliken in 1932.
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 monocentric while a molecular orbital is polycentric.
- The number of molecular orbital formed is equal to the number of combining atomic orbitals. When two atomic orbitals combine, two molecular orbitals are formed. One is known as bonding molecular orbital while the other is called antibonding molecular orbital.
- The bonding molecular orbital has lower energy and hence greater stability than the cor-responding antibonding molecular orbital.
- Just as the electron probability distribution around a nucleus in an atom is given by an atomic orbital, the electron probability distribution around a group of nuclei in a molecule is given by a molecular orbital.
- 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.
Formation of Molecular orbitals
Linear Combination of Atomic Orbitals (LCAO) The atomic orbitals of these atoms may be represented by the wave functions ΨA and ΨB. The formation of molecular orbitals is the linear combination of atomic orbitals that can take place by addition and by subtraction of wave functions of individual atomic orbitals.
The partially positively charged hydrogen atom form a bond with the other more electronegative atom (fluorine, oxygen or nitrogen) is known as hydrogen bond.
Cause of formation of hydrogen bond
- strong electronegative element.
- partial positive and negative charge or polarity of bond or polar molecule.
- Influence of physical state. H2o is a liquid. H2S is a gas.
- High melting point and boiling point.
Types of Hydrogen Bond
- Intermodular Hydrogen bond. It is formed between two different molecules of the same or different compounds.
- Intramolecular Hydrogen bond. It is formed between H atom and highly electronegative atom (F, O, N) present in different bonds with in the same molecule.
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