Chemical Bond : A strong force of attraction holding atoms together in a molecule or crystal by sharing or transfer of electrons.
*Bonds can be strong as ionic bonds,covalent bonds or metallic bonds and weak as hydrogen bonds.
IONIC BONDS : The attractive forces of ionic bond are developed between an electropositive atom and an electronegative atom due to complete transfer of electrons For ex. in NaCl .
Factors Affecting Formation Of Ionic Bond :-
1. Ionisation energy should be low for electropositive elemenet.
2. Electron affinity should be high for electronegative element.
3. Lattice enthalpy should be high.
CHARACTERISTICS :-
1. They generally occur in solid state due to strong electrostatic force of attraction between opposite charged ions.
2. Ionic compounds are generally crystalline in nature.
3. Ionic compounds have generally high melting and boiling points.
4. Ionic compounds are generally soluble in polar solvents and springly soluble in non-polar organic compounds as they follow "like dissolve like".
5. Ionic compounds are good conductors of electricity becouse of presence of mobile ions.
COVALENT BONDS : The bond formed by mutual sharing of electrons between the combining atoms For ex. in Cl2. Sharing of electrons may occur in three ways;
(i) If 2 electrons are shared between two atoms - it forms single bond.
(ii) If 4 electrons are shared between two atoms - it forms double bond.
(iii) If 6 electrons are shared between two atoms - it forms triple bond.
Conditions For Formation Of Covalent Bonds :
1. Electronegativity difference between two atoms should be less than 1.9.
2. When both atoms are short of electrons in valence shell in comparison with noble gas.
COORDINATE BONDS : It is a covalent bond in which the electron pair come from one atom and shared by both the atoms Foe ex. in NH4+.
*(Octet rule has limitations.To overcome, modern theories such as Valence Bond Theory (VBT) and Molecular Orbital Theory (MOT) came.
VALENCE BOND THEORY (ORBITAL CONCEPT) :sproposed by Heitler and London (1927).
Acc. to this theory,a covalent bond is formed by the partial overlap of two half filled atomic orbitals containing electrons with opposite spins.
*Greater is the extent of overlapping ,stronger is bond.
Postulates :
1. This theory deals with the electronoic configuration of the elements.
2. Bond will be formed by overlapping of half filled orbitals.
3. The overlapping atomic orbitals must have comparable energies.
4. Overlapping will increase the electron density and hence stability of molecule.
5. The direction of bond is along the overlapping of atomic orbitals So,covalent bond is directional in nature.
6. covalent bonds are of two types based on the overlapping pattern : (a) sigma bond (b) pi bond.
Limitations : 1. It could not explain the colour,thermodynamics and structural characteristics of compounds.
2. It could not explain the formation of coordinate bonds.
3. It fails to explain the magnetic properties of oxygen and other molecules.
4. This doesn't explain the geometries of some molecules like water,ammonia etc.
Sigma bond : It is a bond which is formed by overlapping of half filled orbitals of two different atoms along their inter-nuclear axis oe end to end overlapping. Sice the overlapping along inter-nuclear axis takes place to large extent Hence sigma bond is strong.It can be formed through s-s overlapping,s-p overlapping and p-p overlapping.
Pi bond : It is formed by lateral or sideways overlapping of p-orbitals i.e. by overlapping of p-orbitals in a direction perpendicular to their inter-nuclear axis.
Valence Shell Electron Pair Repulsion Theory (VSEPR Theory): Proposed by Sidwick and Powell and later on modified by Gillespie and Nyholm.
This theory tells about the shape of molecules on the basis of repulsion of elctron pairs on central atom.
Postulates :
1. Central atom contains two types of electrons :(a) electrons participating in bond formation (bonding pairs) (b) remaining electrons (lone pairs).
2. Negatively charged electron pairs repel each other and try to occupy maximum space to minimise the repulsion.
3. If central atom is linked to similar atoms only and it is surrounded by bond pairs e.g. BF3,CH4,CCl4 then repulsion between bond pairs is si mmilar and molecule has symmetrical shape or regular geometry.
4. If central atom is linked to different atoms or surrounded by bond pairs as well as lone pairs then due to unequal repulsion between electron pairs,molecules aquire unsymmetrical shapes or irregular geometry.
5. The stable geometries of 2,3,4,5,6 bond pairs is linear,triangular,tetrahedral,triangular bipyramidal and octahedral respectively.
6. Lone pair occupies more space on central atom than a bond pair Hence the order of repulsion between electron pairs is :
lone pair - lone pair > lone pair - bond pair > bond pair - bond pair .
Dipolemoment : The polarisation of bonded electron pairs between two atoms is expressed in terms of physical quantity known as dipolemoment.
It is defined as the product of charge (q) and distance (d) between the dis-similar atom i.e. u = qxd
It is a vector quantity which is directed from positive to negative charge. Its unit is Debye(D) and 1 D = 1X10-18 e.s.u.cm
Dipolemoment And Molecular Structure:
1. For Diatomic Molecules, dipolemoment of molecule = dipolemoment of bond.
2. For Polyatomic Molecules, dipolemoment of molecule will be resultant of dipolemoments of all polar bonds.
*Higher is electronegativity difference,greater is dipolemoment.
Applications Of Dipolemoment :
1. In determining polarity of bonds.
2. In calculation of % ionic character.
3. In determination of structure.
4. To distinguish cis and trans isomer.
5. In determining ortho,para and meta configuration.
Hybridisation : The mixing of atomic orbitals of same atoms having slightly different energies so that redistribution of energy takes place between them to form same number of orbitals having identical shapes and equivalent energy. These orbitals are called hybridised orbitals.
Important Points :
1. Only the orbitals of an atom having comparable energies hybridise.
2. Total no. of hybridised orbitals = no. of atomic orbitals.
3. Half filled as well as fully filled orbitals may also take part in hybridisation.
4. Hybridisation never occur in isolated atom but occur when the atom combines with other atom.
5. Type of hybridisation indicates the shape of orbital.
6. The bigger lobe of orbitals are indicated by positive sign and smaller lobe of orbitals are indicated by negative sign.
Types Of Hybridisation :
1. SP Hybridisation: When one s and one p orbital belonging to the same shell of an atom mix together to form two new equivalent orbitals,the type of
hybridisation is known as SP Hybridisation i.e.in BeF2. These two hybrid orbitals are arranged linearly in 3D space with bond angle 180 degree and possess 50 % s and 50 % p character.
2. SP2 Hybridisation: When one s and two p orbitals of the same shell of an atom mix to form three new equivalent orbitals ,the type of hybridisation is SP2 Hybridisation i.e in BF3. These three hybrid orbitals are arranged in 3D space as triangular planar with bond angle 120 degree and possess 33.3% s and 66.6% p character.
3. SP3 Hybridisation: When one s and three p orbitals belonging to the same shell of an atom mix together to form four equivalent orbitals,the type of hybridisation is known as SP3 Hybridisation i.e. in CH4. These four hybrid orbitals are arranged tetrahedrally in 3D spaece with bond angle 109.5 degree and they possess 25% s character and 75% p character.
4. SP3d Hybridisation: When one s,three p and one d orbitals of the same shell of an atom mix to form five new equivalent orbitals ,the type of hybridisation is SP3d Hybridisation i.e in PF5. These five hybrid orbitals are arranged in 3D space as triangular bipyramidal with bond angles 120 degree and 90 degree and possess 20% s,60% p character and 20% d character.
5. SP3d2 Hybridisation: When one s,three p and two d orbitals of the same shell of an atom mix to form six new equivalent orbitals ,the type of hybridisation is SP3d2 Hybridisation i.e in SF6. These five hybrid orbitals are arranged in 3D space as octahedral with bond angle 90 degree and possess 16.6% s,49.8% p character and 33.2% d character.
6. SP3d3 Hybridisation: When one s,three p and three d orbitals of the same shell of an atom mix to form seven new equivalent orbitals ,the type of hybridisation is SP3d3 Hybridisation i.e in IF7. These seven hybrid orbitals are arranged in 3D space as pentagonal bipyramidal with bond angles 72 degree and possess 14.14% s,42.42% p character and 42.42% d character.
Molecular Orbital Theory : proposed by Hund and Mulliken in 1932.
1. Atomic orbitals overlap to form new orbitals ,called Molecular Orbitals.
2. Atomic orbitals are uni-nuclear while molecular irbitals are poly nuclear.
3. As atomic orbitals are energy states of an atom in which electrons of atom are filled,similarly molecular orbitals are energy states of molecules in which electrons of molecules are filled.
4. Only those atomic orbitals combine to form molecular orbital which have comparable energies and proper orientations.
5. The no. of molecular orbitals formed is equal to no. of combining atomic orbitals.
6. Bonding molecular orbital are more stable while Antibonding molecular orbitals are less stable.
7. Filling of electrons in molecular orbitals takes placein accordance with same rules which are followed by atomic orbitals.
8. Shape of molecular orbitals depends upon the type of combining atomic orbitals.
9. Bonding molecular orbitals are represented by sigma,pi and delta whereas Anti-bonding molecular orbitals are represented by sigma star,pi star and delta star.
10. Each electron moving in a molecular orbital has a spin of + 1/2 or -1/2.
Bond order : It is half of difference of bonding and Anti bonding electrons.
B.O. = 1/2 [ Nb - Na ]
where Nb represents the no. of electrons in bonding molecular orbitals and Na is the no. of electrons in anti-bonding molecular orbitals.
Bond length : It is defined as equillibrium distance between centres of nuclie of two bonded atoms.
Factors Affecting Bond Length:
1. Size of atoms : Larger is size of atoms,larger will be the bond length.
2. Multiplicity of bonds : As the multiplicity increases,bond length decreases.
3.Type of hybridisation : SP3 > SP2 SP
Bond energy : For diatomic molecule ,it is the amount of energy to break 1 mole of bond of particular type to separate the atoms.
For poly-atomic molecule ,it is equal to the average of bond dissociation energies of all such bonds.
Factors Affecting Bond Energy:
1. Size of atoms : Smaller is the size of atoms,smaller is bond length ,stronger is bond so larger will be bond energy.
2. Multiplicity of bonds: bond energy is directly proportional to multiplicity of bonds.
3. No. of lone pairs of electron preset : Greater is no. of lone pairs of electrons present in bond,greater is repulsion and less willbe bond energy.
Bond angle : It is the angle between lines representing the direction of bond i.e. in orbitals containing bonding orbitals.
H-Bonding : It is intraction involving a hydrogen atom located between a pair of other atoms having a high affinity for electrons.
Conditions For H - Bonding : 1. The molecule must contain a highly electronegative atom linked to the hydrogen atom.
2. The size of the electronegative atom should be small.
Types of H- bonding : 1. Inter-molecular H-bonding 2. Intra-molecular H-bonding
1. Inter-molecular H-bonding : The hydrogen bonding takes place between different molecules os same or different compounds,is called inter-molecular H-bonding For e.g in water,ammonia.
2. Intra-molecular H-bonding : The hydrogen bonding takes place with in a molecule itself is called intra-molecular H-bonding For e.g. in Ortho-nitrophenol.
Effects Of H- Bonding:
1. Association : Molecules get dimerised,trimerised or polumerised like carboxylic acids get dimerised and molecular mass is found to be double.
2. High melting and boiling points : It is found that the compounds whose molecules have inter- molecular H-bonding with each other have high melting and boiling points.
3. Solubility : The molecules containing H-bonding have more solubity in water like alcohols.