CHEMICAL BONDING AND MOLECULAR STRUCTURES

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Chemical Bonding Part-1

Chemical Compounds

  1. Ionic Compounds
  2. Covalent Compounds
  3. Co-Ordinate Compounds

Chemical Bond

  1. Ionic Bond
  2. Covalent Bond
  3. Co-ordinate Bond
  4. Metallic Bond
  5. Hydrogen Bond
  6. Vandaar Wall Bond

Compound

  • When two or more elements combine in a fixed proportion then the species developed is known as compound.

Chemical Bond

  1. During the compound formation, the elemental interaction developed is known as chemical bond.
  2. The elements go for compound formation to achieve stability from its isolated state.
  3. In general the Inert Gas Configuration is the most stable arrangement. So through compound formation elements try to achieve inert gas configuration.
  4. In inert gas configuration the element achieves duet or octet configuration. These are most stable configuration.
  5. In compound formation there is involvement of only valence electrons.

Lewis-Kossel Octet Rule

  1. The species which have the complete octet will be more stable.
  2. Elements try to achieve octet configuration by exchange of electrons.
  3. To achieve the octet configuration elements may release the electrons or may gain electrons. Due to this formation of ions take place.
  4. To achieve the octet configuration elements may share their electrons.

Lewis-Electron  Structure

  • To represent the electronic structure there in involvement of only valence electrons.

Covalent bond

  1. When equal sharing of electrons take place then the bond developed is known as covalent bond.
  2. By the sharing of two electrons in between two centers one covalent bond is developed.
  3. Bond Pair : Pair of electrons individually represent on a particular center in a molecule.
  4. Lone Pair : Pair of electrons in between two centers in a molecule.

Hypervalent Molecules

  1. The molecule in which expansion of octet takes place are known as hypervalent molecule.
  2. In hypervalent molecule the individual center which has more than 8 electrons is known as hypervalent center.

Hypovalent Molecules

  1. The molecule in which octet is not completed is known as hypovalent molecule.
  2. The individual center which has less than 8 electrons is known as hypovalent center.

Limitations of Octet Rule

  • As per the octet rule only the species which have complete octet would be stable, but there are number of molecules which do not follow the octet rule but they are stable.

Expansion of Octet

  • There are number of molecules which have more than 8 electrons at a particular center and are stable.

Odd electrons molecule

  • The molecules which have the odd electrons in their valence shell cannot be explained by Lewis octet rule.

Note

  1. Compounds of the inert gas element as per the Lewis octet rule the element with inert gas configuration is most stable and should not participate in the reaction.
  2. But the inert gas element itself makes the number of compound and participate in the reaction.

Formal Charge

  • It is an imaginary charge which appears on an element by considering no difference in their electronegativity.

Resonance

  1. When properties of a given species cannot be explained only by one structure then various structures have to be written.
  2. There structures are called resonating forms or canonical form.
  3. By considering all the canonical form one average structure is written known as resonance hybrid. This phenomena is known as resonance.
  4. Resonance is an imaginary concept, there is no existence of canonical forms but resonance hybrid is true structure.

Covalent Molecule

  • Covalent compounds are the compounds which are formed by sharing of electrons.

Theories for Molecule

Valence Bond Theory

  1. VBT was proposed by the scientist Heitler and London. It is based on the concept of overlapping.
  2. When two elements are isolated from each other they have their own particular energy.
  3. From elemental state to compound state the interaction takes place between two elements. There is attractive as well as the repulsive force.
  4. The attractive forces operate between nuclear part of one element to the electrons of the other element. Similarly repulsive forces operate between nuclear part as well as electron cloud of both elements.
  5. During process of the formation of compound the element comes closure to each other, in this attractive forces dominate over the repulsive forces. Here the energy of the system decreases.
  6. At a particular point the system achieves the minimum energy. Here attractive forces counter balance the repulsive forces and bond formation takes place.
  7. At this stage the distance between the two centers is the bond length and the energy is known as bond energy. Here the system is most stabilized.
  8. On bringing the center nearer to this the repulsive forces dominate over the attractive forces and system becomes unstable and the energy of the system increases. Hence there will be no existence of the molecule.

Overlapping

  1. In the overlapping there is involvement of the valence shell atomic orbitals.
  2. The atomic orbitals which are full filled do not participate in overlapping
  3. Only those atomic orbitals which have unpaired electrons can participate in overlapping.
  4. To do the overlapping of the atomic orbitals initially axis of bond formation has to be decided i.e. internuclear axis.
  5. As the orbitals approach from the particular axis, hence covalent bond is directional.
  6. The strength of the bond depends upon the extent of overlapping, greater overlapping stronger would be the bond.

Types of Overlapping

  1. Head-on Overlapping
  2. Lateral Overlapping
  3. Zero Overlapping

Head-on interaction

  1. A p-orbital can overlap with a p-orbital.
  2. As the p-orbitals are perpendicular to each other hence a Px-orbital can overlap with the Px-orbital.
  3. The overlapping of orbitals can be constructive or destructive.
  4. In a constructive phenomena the waves combines with each other and this is positive overlapping.
  5. In destructive phenomena the waves cancel each other and this is negative overlapping.
  6. For bond formation there should be positive overlapping.
  7. When bond formation takes place directly on the inter nuclear axis, then it is head to head, head-on interaction. The bond developed through this interaction is know is sigma bond.

Lateral Overlapping

  1. Overlapping of the orbitals can take place above and below the inter nuclear axis, then this is known as lateral overlapping or side ways overlapping.
  2. For example if z-axis is decided as the axis of bond formation then on the inter nuclear axis there is overlapping of Pz-Pz and formation of sigma bond.
  3. From the z-axis the perpendicular orbitals are Px and Py.
  4. If combination of p-orbitals take place above and below the inter nuclear axis then this is lateral overlapping and formation of pi-bond.

Sigma Bond

  1. This is the bond developed on the inter nuclear axis through head to head overlapping.
  2. Here, extent of overlapping is high, hence sigma bond would be stronger.
  3. Between two centers only one sigma bond can be formed.

Pi Bond

  1. This is the bond developed above and below the inter nuclear axis, through lateral overlapping.
  2. Here extent of overlapping is less, hence it is weak bond with respect to individual sigma bond.
  3. Between two centers more than one pi bond could be formed.

Note

  1. As per VBT oxygen should be diamagnetic but experimentally it is found paramagnetic.
  2. VBT could not explain this property of the molecules.
  3. Delta Bond : When all the four lobes of d-orbital participate in the overlapping then it is delta bond.

CHEMICAL BONDING AND MOLECULAR STRUCTURES PART-2

VSEPR Theory

  1. It was proposed by Gillipse and Nyholm.
  2. Defined the geometry of the molecules.
  3. As per the VSEPR theory each molecule have one particular three dimensional arrangement.
  4. This particular arrangement is most stabilized and least energetic.
  5. The molecule achieve this stable arrangement by three dimensional arrangement of the element in such a way that there should be minimum electronic repulsion.
  6. A lone pair is accumulated on an individual center and would be exposed more in the space.
  7. A bond pair is accumulated between two centers and its electron cloud would be concentrated in small space. It would be under the control of two nucleus.
  8. According to this electronic repulsion order would be: lp-lp > lp-bp > bp-bp.
  9. Due to any reason if molecule deviates form its regular arrangement then its geometry gets alter and new shape is formed.

Hybridization

  1. Hybridization concept was given by Pauiling.
  2. In hybridization process there is mixing of two or more different types of atomic orbitals.
  3. In hybridization there is involvement of atomic orbitals, hence fulfilled, half-filled and empty atomic orbitals participate.
  4. In hybridization excitation of electrons is allowed if tis is required.
  5. In hybridization there is exchange of energy among different atomic orbitals. Hence, energy gets exchanged and atomic orbital gets intermixed. Then there is formation of new type of orbitals known as hybrid orbitals.
  6. The hybrid orbitals have similar geometry and equal energy i.e. they are degenerate orbitals.
  7. Number of hybrid orbitals would be equal toe the number of atomic orbitals intermixed.
  8. Hybrid orbitals are directional in nature and develop only the sigma bond.
  9. Hybrid orbitals do not form the pi-bonds.
  10. The hybrid orbitals arrange themselves in such a way that there should be minimum electronic repulsion.
  11. Hybrid orbital can overlap with any other atomic orbital to form the sigma bond.
  12. Hybridization is represented by the number of atomic orbitals.

    No. of hybridization O’s = No. of atomic O’s intermixed

HybridizationTypeGeometry ShapeBond Angle
SpAB2linearlinear180
Sp2AB3Trigonal planarTrigonal planar120
AB2LTrigonal planarBent<120
Sp3AB4TetrahedralTetrahedral109
AB3LTetrahedralPyramidal107
AB2L2TetrahedralBent104
Sp3dAB5Trigonal bipyramidalTrigonal bipyramidal90, 180, 120
AB4LTrigonal bipyramidalSee-saw
AB3L2Trigonal bipyramidalT-shape
AB2L3Trigonal bipyramidalLinear
Sp3d2AB5OctahedralOctahedral900
AB4L1OctahedralSquare pyramidal
AB3L2OctahedralSquare planar

Sp3d

  1. Sp3d hybridization is a combination of sp2 and pd hybrid orbitals.
  2. Its regular arrangement is Trigonal bi-pyramidal in which sp2 orbital forms the Trigonal arrangement where as pd orbital forms the pyramidal arrangement.
  3. The Trigonal arrangement positions are known as equatorial position where as pyramidal positions are known as axial position.
  4. In general axial bond length is higher than equatorial bond length.
  5. It is based on electronic repulsions.

Bent Rule

  • In sp3d hybridization a more electronegative element occupy the axial position whereas a lone pair occupies the equatorial position.

Note

  1. In PF5 axial and equatorial bond lengths are equal whereas in other phosphorus pentahalides axial and equatorial bond lengths are different.
  2. Reason : In PF5 there is rapid exchange of the chlorine atom on equatorial and axial position and hence all the bond length becomes similar and this is known as pseudo rotation.
  3. PI5 does not exist on one phosphorus. Five iodine cannot be stabilized due to steric crowding.
  4. PCl5 exist but NCl5 doesn’t exist.
  5. Reason : In nitrogen there is no d-orbital hence it cannot expand the octet (sp3d not possible).

Sp3d2

  1. In Sp3d2 hybridization for AB6 molecules the geometry is octahedral.
  2. All the six positions are identical and 900 to each other.
  3. In case of AB5L molecule that is when one lone pair is present then it can occupy any of the positions as all the positions are identical.
  4. In case of AB4L2 molecules when two lone pairs are present then lone pair should be opposite to each other to minimize the electronic repulsion.
  5. Here lone pair cannot be on the adjacent position.

Chemical Bonding Part-3

Dipole moment

  1. Dipole moment define the flow of electron density from a less electronegative terminal to more electronegative terminal.
  2. It is a vector property.
  3. Unit for dipole moment is debye.
  4. The molecule which have the dipole moment would be polar in nature where the molecule in which ne dipole becomes zero would be non-polar.img_chemical bonding_01

Hydrogen bond

  1. When hydrogen comes in between the two electronegative elements like nitrogen, oxygen and fluorine and from one electronegative element it is bonded with the covalent bond whereas, from other electronegative element it develop the ionic interaction.
  2. This ionic interaction is known as hydrogen bond.
  3. Hydrogen bond is not a true bond as there is neither transfer nor sharing of electron density.
  4. Hydrogen bond is a weak bond.

Types of hydrogen bond

  1. Inter molecular hydrogen bond.
  2. Intra molecular hydrogen bond.

Inter molecular hydrogen bond

  1. When hydrogen bonding takes place in between the molecules then it is intermolecular hydrogen bond.
  2. The molecules may be similar type or different type.
  3. Due to intermolecular hydrogen bonding number of molecule gets associated with each other.

Intra molecular hydrogen bond

  1. The molecules in which hydrogen bonding takes place within the molecule then it is intra molecular hydrogen bond.
  2. In case of intra molecular hydrogen bonding there in no association of any other molecule.

Examples of intra molecular hydrogen bonding

  • Ortho-distributed benzene
  • In case of ortho-distributed benzene whichever group like: OH,NO2,NH3 etc will develop intra molecular hydrogen bond.
  • Through intra molecular hydrogen bonding there is ring formation also known as chelation.
  • Acetylacetone
  • Choraldihydrate
  • Nickel dimethyle glyoxime
  • Hydrated Cu2SO4

Effects of hydrogen bonding

1. Solubility of organic compounds in water

  1. Alcohols are soluble in water but alkylhalides or alkanes are insoluble in water.
  2. This is due to that alcohol can easily develop the hydrogen bonding with water but alkylhalides or alkane cannot.
  3. Any compound which can develop hydrogen bonding with the water is soluble in water.
  4. Lower molecular mass alcohol are water soluble but higher molecular mass alcohol are water insoluble.
  5. In case of the higher molecule mass alcohol the bulky alkyl group would be hydrophobic in nature and it does not allow the hydrogen bonding.

2.Volatile Behavior of Compound or Evaporation tendency

  1. Ethers are more volatile than alcohol and similarly acetones are more volatile than alcohol.
  2. In alcohol due to hydrogen bonding the molecules get associated with each other, hence the rate of evaporation decreases. Whereas in acetone and ethers there is no hydrogen bonding hence evaporation is faster.
  3. Volatile behavior: The molecules which can easily make the vapor easily that is has the more evaporation tendency are more volatile compound.
  4. More volatile compound would be having less intermolecular forces.

  • The molecules which have the more hydrogen bonding would be more viscous as they have more association.

  • Compounds which have the more hydrogen bonging would have higher boiling point as they require more energy to go into the gaseous phase.

  1. Density of the ice is less than water due to this ice floats on water.
  2. In ice the water molecule arrange tetrahedral and each water molecule makes the four hydrogen bond.
  3. In his case there is empty space which gets developed and ice has the cage like structure.
  4. Due to this cage structure volume of the ice would be higher and density would be less.

Vander Waals Forces

Ionic interaction

  1. In ionic interaction there is a positive-negative interaction.
  2. In this case a cationic terminal gets surrounded by anionic terminal and an anionic terminal gets surrounded by cationic terminal.
  3. This is one of the strongest interaction.

Ion-dipole interaction

  1. In ion-dipole interaction one of the molecule should be ionic compound whereas the other molecule should be polar covalent compound.
  2. Here as the two compounds kept together then cationic terminal would get surrounded by the negative end of the dipole whereas an anionic terminal would get surrounded by the positive terminal of the dipole.
  3. This interaction is known as ionic dipole interaction.

Dipole-Dipole interaction

  1. Dipole-dipole interaction is also known as keesom forces.
  2. In dipole-dipole interaction there are two polar covalent molecule.
  3. As the molecules interact then they orient themselves in such a way that positive end of one dipole molecule would be towards the negative end of the other dipole molecule and vice-versa.

Dipole-Induced Dipole

  1. In dipole-induced dipole interaction there is one polar molecule and the other is non-polar molecule.
  2. When these types of the molecules interact in non-polar molecule partially polarity gets developed in presence of polar molecule.
  3. This dipole-induced dipole interaction also known as debye forces.

Induced Dipole-Induced Dipole

  1. In induced dipole-induced dipole interaction both the molecules are non-polar.
  2. This is known as London forces or dispersion forces.
  3. This is the weakest interaction.
  4. In induced dipole-induced dipole interaction when non-polar molecule interact then partially dipole gets induced.

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