Briefly describe the valence bond theory of covalent bond formation by taking an example of hydrogen. How can you interpret energy changes taking place in the formation of dihydrogen?
The valence bond theory was put forward by Heitler and London in 1927. It was later improved and developed by L. Pauling and J.C. Slater in 1931. The valence bond theory is based on the knowledge of atomic orbitals and electronic configurations of elements, overlap criteria of atomic orbitals and stability of molecule.
The main points of valence bond theory are
(i) Atoms do not lose their identity even after the formation of the molecule.
(ii) The bond is formed due to the interaction of only the valence electrons as the two atoms come close to each other. The inner electrons do not participate in the bond formation.
(iii) During the formation of bond, only the valence electrons from each bonded atom lose their identity. The other electrons remain unaffected.
(iv) The stability of bond is accounted by the fact that the formation of bond is accompanied by release of energy. The molecule has minimum energy at a certain distance between the atoms known as intemuclear distance. Larger the decrease in energy, stronger will be the bond formed.
Valence bond Treatment of Hydrogen Molecule:
Consider two hydrogen atoms A and B approaching each other havingnuclei Ha and HB and the corresponding electrons eA and eB respectively.
When atoms come closer to form molecules new forces begin to operate.
(a) The force of attraction between nucleus of atom and electron of another atom.
(b) The force of repulsion between two nuclei of the atom and electron of two atoms.
Fig. (a) Two hydrogen atoms at a large distance and hence, no interaction, (b) Two hydrogen atom closer to each other atomic orbitals begin to interact, (c) Attractive and repulsive forces in hydrogen atoms when interaction begins. In case of hydrogen: Figure ‘a' shows that two hydrogen atoms are at farthest distances and their electron distribution is absolutely symmetrical.
(a) When two hydrogen atom start coming closer to each other, the electron cloud becomes distorted and new attractive and repulsive forces begin to operate as shown in figure ‘c'
(b) In figure ‘c' dotted lines show attractive forces present in atom already and bold lines show the new attractive and repulsive forces.
(c) It has been found experimentally that the magnitude of net attractive forces is more than net repulsive forces. Thus stable hydrogen molecule is formed.
Potential energy diagram for formation of hydrogen molecules:
When two hydrogen atoms are at farther distance, there is no force operating between them, when they start coming closer to each other, force of attraction comes into play and their potential energy starts decreasing. As they come closer to each other potential goes on decreasing, but a point is reached, when potential energy acquires minimum value.
Note:
(a) This distance corresponding to this minimum energy value is called the distance of maximum possible approach, i.e. the point which corresponds to minimum energy and maximum stability.
(b) If atoms come further closer than this distance of maximum possible approach, then potential energy starts increasing and force of repulsion comes into play and molecules starts becoming unstable.
Elements X, Y and Z have 4, 5 and 7 valence electrons respectively, (i) Write the molecular formula of the compounds formed by these elements individually with hydrogen, (ii) Which of these compounds will have the highest dipole moment?
Which of the following statements are not correct?
(a) NaCl being an ionic compound is a good conductor of electricity in the solid state.
(b) In canonical structures there is a difference in the arrangement of atoms.
(c) Hybrid orbitals form stronger bonds than pure orbitals.
(d) VSEPR theory can explain the square planar geometry of XeF4.
What is the effect of the following processes on the bond order in N-, and 02?
(i) N2 → N+2 + e– (ii) 02 → O+2 + e–
Briefly describe the valence bond theory of covalent bond formation by taking an example of hydrogen. How can you interpret energy changes taking place in the formation of dihydrogen?
Match the items given in Column I with examples given in Column II.
| Column I | Column II |
| (i) Hydrogen bond | (a) C |
| (ii) Resonance | (b) LiF |
| (iii) Ionic solid | (c) H2 |
| (iv) Covalent solid | (d) HF |
| (e) 03 |
Structures of molecules of two compounds are given below:
(a) Which of the two compounds will have intermolccular hydrogen bonding and which compound is expected to show intramolecular hydrogen bonding?
(b) The melting point of a compound depends on. among other things, the extent of hydrogen bonding. On this basis explain which of the above two compounds will show higher melting point.
(c) Solubility of compounds in water depends on power to form hydrogen bonds with water. Which of the above compounds will form hydrogen bond with water easily and be more soluble in it?
Draw diagrams showing the formation of a double bond and a triple bond between carbon atoms in C2 H4 and C2 H2 molecules.
Write Lewis structure of the following compounds and show formal charge on each atom. HN03, No2, H2so4
Assertion (A): Though the central atom of both NH3 and H20 molecules are sp3 hybridised, yet H – N – H bond angle is greater than that of H – O – H.
Reason (R): This is because nitrogen atom has one lone pair and oxygen atom has two lone pairs.
(a) A and R both are correct, and R is the correct explanation of A.
(b) A and R both are correct, but R is not the correct explanation of A.
(c) A is true but R is false.
(d) A and R both are false.
Although both CO2 and H2O are triatomic molecules, the shape of H2O molecule is bent while that of CO2 is linear. Explain this on the basis of dipole moment.
Is there any change in the hybridisation ofB and N atoms as a result of the following reaction ? BF3 + NH3 ——-> F3 B.NH3
In which of the following molecule/ion all the bonds are not equal?
(a) XeF4
(b) BF–4
(c) C2H4
(d) SiF4
Diamagnetic species are those which contain no unpaired electrons. Which among the following are diamagnetic?
(a) N2
(b) N22-
(c) 02
(d) o22-
Explain the non linear shape of H2S and non planar shape of PCl3 using valence shell electron pair repulsion theory.
Using molecular orbital theory, compare the bond energy and magnetic character of 0+2 and O–2
What is the total number of sigma and pi bonds in the following molecules?
(a) C2 H2 (b) C2 H4
Compare the relative stability of the following species and indicate their magnetic properties: O2, O2, O2– (Superoxide),O22- (peroxide)
Which molecule/ion out of the following does not contain unpaired electrons?
(a) N+2
(b) 02
(c) O22-
(d) B2
3PO3 can be represented by structures 1 and 2 shown below. Can these two structures be taken as the canonical forms of the resonance hybrid representing H3PO3? If not, give reasons for the same.
Arrange the bonds in order of increasing ionic character in the molecules: LiF, K2O, N2, SO2 and ClF3.
Explain why BeH2 molecule has a zero dipole moment although the Be—H bonds are polar.
Describe the change in hybridisation (if any) of the Al atom in the following reaction. AlCl3 + Cl– ——>AlCl4- .
Write the important conditions required for the linear combination of atomic orbitals to form molecular orbitals.
Define Lattice energy. How is Lattice energy influenced by (i) Charge on the ions (ii) Size of the ions?
Give reasons for the following: ‘
(a) Covalent bonds are directional bonds while ionic bonds are non- directional.
(b) Water molecule has bent structure whereas carbon dioxide molecule is linear.
(c) Ethyne molecule is linear.
Match the species in Column I with the type of hybrid orbitals in Column II.
| Column I | Column II |
| (i) SF4 | (a) sp3cf |
| (ii) if5 | (b) d2sp3 |
| (iii) NO2+ | (c) sp3 d |
| (iv) NH4 | (d) sp3 |
| (e) sp |
Match the species in Column I with the geometry/shape in Column II.
| Column I | Column II |
| (i) H30+ | (a) Linear |
| (ii) HC = CH | (b) Angular |
| (iii) Cl0–2 | (c) Tetrahedral |
| (iv) NH+4 | (d) Trigonal bipyramidal |
| – | (e) Pyramidal |
Match the shape of molecules in Column I with the type of hybridization in Column II.
| Column I | Column II |
| (i) Tetrahedral | (a) sp2 |
| (ii) Trigonal | (b) sp |
| (iii) Linear | (c) sp3 |
Write Lewis symbols for the following atoms and ions: S and S2– ; Al and Al3+; H and H–
Draw the Lewis structures for the following molecules and ions:
H2S, SiCl4 , BeF2, C032-, HCOOH
