Chemistry

Chemical Bonding and Molecular Structure

Question:

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?

Answer:

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.

ncert-exemplar-problems-class-11-chemistry-chapter-4-chemical-bonding-and-molecular-structure-63
ncert-exemplar-problems-class-11-chemistry-chapter-4-chemical-bonding-and-molecular-structure-64
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.

ncert-exemplar-problems-class-11-chemistry-chapter-4-chemical-bonding-and-molecular-structure-65

previuos
next

Chemical Bonding and Molecular Structure

Q 1.

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?

Q 2.

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.

Q 3.

State the types of hybrid orbitals associated with (i) P in PCl5  and (ii) S in  SF6

Q 4.

Write the significance of plus and minus sign in representing the orbitals,

Q 5.

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

Q 6.

Which of the following species have the same shape?
(a) C02
(b) CC14                                  
(c) 03                                                
(d) N02

Q 7.

Write Lewis structure of the following compounds and show formal charge on each atom.  HN03, No2, H2so4

Q 8.

All the C – O bonds in carbonate ion (CO2-3) are equal in length. Explain.

Q 9.

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?

Q 10.

Define electronegativity. How does it differ from electron gain enthalpy?

Q 11.

Name the two conditions which must be satisfied for hydrogen bonding to take place in a molecule.

Q 12.

Species having same bond order are
(a) N2                                            
(b) N2                                              
(C) F+2                                            
(d) o2

Q 13.

Explain why PC15 is trigonal bipyramidal whereas IF5 is square pyramidal.

Q 14.

What is an ionic bond? With two suitable examples explain the difference between an ionic and covalent bond?

Q 15.

Arrange the following bonds ‘in order of increasing ionic character giving reason.
N-H, F-H, C-H and O-H

Q 16.

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.

Q 17.

Arrange the bonds in order of increasing ionic character in the molecules: LiF, K2O, N2, SO2 and ClF3.

Q 18.

Is there any change in the hybridisation ofB and N atoms as a result of the following reaction ?  BF3 + NH3 ——-> F3 B.NH3

Q 19.

Considering X-axis as the intemuclear axis which out of the following will not form a sigma bond and why? (a) Is and Is (b) Is and  2px  (c)  2py  and 2py (d) Is and 2s

Q 20.

Predict the shapes of the following molecules using VSEPR theory?
(i) BeCl2(ii) SiCl4

Q 21.

Which is more polar CO2  or  N2O? Give reason.

Q 22.

(a) Define dipole moment. What are the units of dipole moment?
(b) Dipole moment values help in predicting the shapes of covalent molecules. Explain.

Q 23.

(a) How many a and n bonds are present in
ncert-solutions-for-class-11-chemistry-chapter-4-chemical-bonding-and-molecular-structure-34
(b) Why Hf is more stable than H2?
(c) Why is B2 molecule paramagnetic?

Q 24.

Which of the following statements are correct about CO32- ?
(a) The hybridization of central atom is sp3.
(b) Its resonance structure has one C – O single bond and two C = O double bonds.
(c) The average formal charge on each oxygen atom is 0.67 units.
(d) All C – O bond lengths are equal.

Q 25.

Structures of molecules of two compounds are given below:

ncert-exemplar-problems-class-11-chemistry-chapter-4-chemical-bonding-and-molecular-structure-26

(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?

Q 26.

Assertion (A): Among the two O – H bonds in H20 molecule, the energy required to break the first O – H bond and other O – H bond is the same.
Reason (R): This is because the electronic environment around oxygen is the same even after breakage of one O – H bond.
(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.

Q 27.

Write Lewis symbols for the following atoms and ions: S and  S2– ; Al and  Al3+; H and H

Q 28.

Apart from tetrahedral geometry, another possible geometry for CH4 is square planar with the four H atoms at the comers of the square and the C atom at its centre. Explain why CH4 is not square planar?

Q 29.

Which out of NH3 and NF3 has higher dipole moment and why?

Q 30.

Explain the formation of  H2 molecule on the basis of valence bond theory.

Q 31.

Define hydrogen bonds. Is it weaker or stronger than the van der Waals forces?

Q 32.

Define covalent bond according to orbital concept?

Q 33.

Out of bonding and antibonding molecular orbitals, which one has lower energy and which one has higher stability?

Q 34.

In which of the following molecule/ion all the bonds are not equal?
(a) XeF4                                  
(b) BF4                                        
(c) C2H4                                    
(d) SiF4

Q 35.

Which of the following attain the linear structure?
(a) BeCl2
(b) NCO+                                    
(c) N02                                          
(d) CS2

Q 36.

CO is isoelectronic with
(a) NO+
(b) N2                                              
(c) SnCl2                                    
(d) N02

Q 37.

Q 38.

Explain why CO2-3 ion cannot be represented by a single Lewis structure. How can it be best represented?

Q 39.

Draw the resonating structure of (i) Ozone molecule (ii) Nitrate ion

Q 40.

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

Q 41.

Explain the formation of a chemical bond.

Q 42.

How do you express the bond strength in terms of bond order?

Q 43.

Use Lewis symbols to show electron transfer between the following atoms to form cations and anions (a) K and S (b) Ca and O (c) Al and N.

Q 44.

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.

Q 45.

The skeletal structure of  CH3COOH  as shown below is correct, but some of the bonds are shown incorrectly. Write the correct Lewis structure for acetic acid.
ncert-solutions-for-class-11-chemistry-chapter-4-chemical-bonding-and-molecular-structure-12

Q 46.

What is the total number of sigma and pi bonds in the following molecules?
(a) C2 H2 (b) C2 H4

Q 47.

Write the important conditions required for the linear combination of atomic orbitals to form molecular orbitals.

Q 48.

Compare the relative stability of the following species and indicate their magnetic properties: O2, O2, O2 (Superoxide),O22- (peroxide)

Q 49.

Out of sigma and Π  bonds, which one is stronger and why?

Q 50.

Arrange  O2,O2,O22-, O2+in increasing order of bond energy.