Question

3. Write the energies of molecular orbital of species upto \[14\] electrons and after \[14\] electrons
4. Calculate the bond order of the following species: \[{{N}_{2}},{{O}_{2}},{{O}_{2}}^{2-},{{N}^{2+}},N{{e}_{2}},{{F}_{2}},B{{e}_{2}}\]
5. Define hydrogen bonding. Types of hydrogen bonding with examples.

Answer 1

Hint:
-The filling of electrons in the molecular orbital is based on the theory of molecular orbitals, according to which the atomic orbitals combine in order to form the molecular orbitals of the same energy.
-The presence of hydrogen bonding in a molecule makes it more stable, because the hydrogen involved in the process experiences much more attractive forces, which prevents the dissociation of proton from the molecule under consideration.

Complete answer:
In molecular orbital theory, electrons in a molecule are not assigned to the chemical bonds between individual atoms, but are observed as moving under the influence of the atomic nuclei in the whole molecule. This theory takes into account the existence of bonding and antibonding molecular orbitals. Bonding molecular orbitals are those orbitals which are formed by in-phase combinations of atomic wave functions, and the electrons present in these orbitals stabilize a molecule, or the molecule is more stable in this case. Whereas, antibonding molecular orbitals are a result of out-of-phase combinations of atomic wave functions and electrons which are present in these orbitals make a molecule less stable.
Energy of Bonding Molecular Orbitals is less than the energy of Anti Bonding Molecular Orbitals due to the attraction of both the nuclei for both the electrons is higher.
Energy of Anti Bonding Molecular Orbitals is higher than Bonding Molecular Orbitals as the electrons try to move away from the nuclei and are in a repulsive state.
\[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{2}}{{\left( 2\pi \right)}^{4}}{{\left( 2\sigma \right)}^{2}}{{\left( 2\pi * \right)}^{4}}\]
Where \[\sigma \]represents the bonding molecular orbitals and \[\sigma *\] represent the antibonding molecular orbitals.
4. Bond order is simply the number of covalent bonds present in a covalent molecule. It is equal to one half of the difference between the numbers of electrons in the antibonding & bonding molecular orbitals. The general formula for calculation of bond order is
\[Bond\text{ }order\text{ }=\text{ }\dfrac{\left( {{N}_{b}}-{{N}_{a}} \right)}{\text{ }2}\]
Where, \[{{N}_{a}}\] denotes the number of electrons in the antibonding orbitals and \[{{N}_{b}}\]represents the number of electrons present on the bonding molecular orbital.
In case of \[{{N}_{2}}\] the bond order of can be calculated as follows:
Electronic configuration: \[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{2}}{{\left( 2\pi \right)}^{4}}{{\left( 2\sigma \right)}^{2}}\]
Here, \[{{N}_{b~}}=\text{ }8\] and \[{{N}_{a}}~=\text{ }2\]
\[B.O\text{ }=\text{ }\dfrac{\left( 8-2 \right)}{2}\]
\[B.O\text{ }=\text{ }3\]
In case of\[{{O}_{2}}\] the bond order of can be calculated as follows:
Electronic configuration:
\[B.O\text{ }=\text{ }\dfrac{\left( 10-6 \right)}{2}\]
\[B.O\text{ }=2\]
In case of\[{{O}_{2}}^{2-}\] the bond order of can be calculated as follows:
Electronic configuration: \[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{2}}{{\left( 2\pi \right)}^{4}}{{\left( 2\sigma \right)}^{2}}{{\left( 2\pi * \right)}^{3}}\]
\[B.O=1/2\text{ }\left( 8-5 \right)\text{ }=\text{ }1.5\]
In case of\[{{N}_{2}}^{+}\] the bond order of can be calculated as follows:
Electronic configuration: \[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{1}}{{\left( 2\pi \right)}^{4}}{{\left( 2\sigma \right)}^{1}}\]
\[B.O=\text{ }\left( 94 \right)/2\text{ }=\text{ }5/2\text{ }=\text{ }2.5\]
In case of\[N{{e}_{2}}\] the bond order of can be calculated as follows:
Electronic configuration: \[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{2}}{{\left( 2\pi \right)}^{4}}{{\left( 2\sigma \right)}^{2}}{{\left( 2\pi * \right)}^{4}}{{\left( 2\sigma * \right)}^{2}}\]
Bond order is \[{\scriptscriptstyle 1\!/\!{ }_2}\text{ }\left( 8\text{ }\text{ }8 \right)\text{ }=\text{ }1\]
In case of\[{{F}_{2}}\] the bond order of can be calculated as follows:
Electronic configuration: \[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{2}}{{\left( 2\pi \right)}^{4}}{{\left( 2\sigma \right)}^{2}}{{\left( 2\pi * \right)}^{4}}\]
Bond order is \[{\scriptscriptstyle 1\!/\!{ }_2}\text{ }\left( 8\text{ }\text{ }6 \right)\text{ }=\text{ }1\]
In case of\[B{{e}_{2}}\] the bond order of can be calculated as follows:
Electronic configuration: \[{{\left( 1\sigma \right)}^{2}}{{\left( 1\sigma * \right)}^{2}}\]
Bond order is \[{\scriptscriptstyle 1\!/\!{ }_2}\text{ }\left( 2\text{ }\text{ }2 \right)=0\]
5. Hydrogen bonding is a special type of bonding which involves dipole-dipole attraction between molecules, and not a covalent bond to a hydrogen atom. It is a result of the attractive force present between a hydrogen atom which is covalently bonded to a very electronegative atom such as a \[N,\text{ }O,\text{ }or\text{ }F\] atom and another very electronegative atom.
There are two types of Hydrogen bonding, intermolecular and intramolecular.
When hydrogen bonding occurs between different molecules of the different or same compounds, it is called intermolecular hydrogen bonding. For instance hydrogen bonding present in water, alcohol, ammonia etc.
On the other hand the hydrogen bonding which occurs within a molecule itself, is termed as intramolecular hydrogen bonding. For instance, the hydrogen bonding present in acetone, etc.

Note: Two very common properties of hydrogen bonding are,
-Solubility: Lower alcohols as in alcohols with lesser number of carbons, are soluble in water because of the existence of hydrogen bonding which can take place between water and alcohol molecules.
-Volatility: Because the compounds involving hydrogen bonding between different molecules have much higher boiling point, this is why they are less volatile.