Chemical Bonding and Molecular Structure

Why do atoms combine, and why do molecules have definite shapes? This chapter begins with the Kossel-Lewis approach, in which atoms achieve a stable octet by transferring electrons (ionic bonds) or sharing them (covalent bonds), using Lewis dot structures and formal charges. The octet rule has limitations: incomplete octets in BCl3, expanded octets in PCl5 and SF6, and odd-electron molecules like NO. You will study ionic bond formation and lattice enthalpy (788 kJ/mol for NaCl). Bond parameters include bond length, bond angle, bond enthalpy and bond order. Resonance in O3 and CO3 2- and bond polarity (dipole moment) are also covered. The VSEPR theory explains molecular shapes from the repulsions between bond pairs and lone pairs, giving shapes such as linear BeCl2, trigonal planar BF3, tetrahedral CH4, pyramidal NH3 (107 degrees) and bent H2O (104.5 degrees). Valence bond theory explains bonds through orbital overlap, sigma and pi bonds, and hybridisation (sp, sp2, sp3, sp3d and sp3d2). Molecular orbital theory explains bond order and the paramagnetism of O2. Finally, hydrogen bonding is discussed.

Detailed Worksheet: Chemical Bonding and Molecular Structure Section A - Definitions (10 marks) 1. What is the octet rule? State two of its limitations. (2 marks) 2. What is lattice enthalpy? (2 marks) 3. Define bond order. Write the bond order of N2 and O2. (2 marks) 4. Distinguish between sigma and pi bonds. (2 marks) 5. What is hydrogen bonding? Give one example each of intermolecular and intramolecular hydrogen bonding. (2 marks) Section B - Calculations and Applications (15 marks) 6. The bond length of HCl is 127 pm and its dipole moment is 1.03 D. Calculate the dipole moment if the bond were 100% ionic (charge = 1.602 x 10^-19 C; 1 D = 3.336 x 10^-30 C m) and hence the percentage ionic character of the bond. (3 marks) 7. Calculate the formal charges on the three oxygen atoms in the Lewis structure of O3, using formal charge = valence electrons - non-bonding electrons - (bonding electrons)/2. (3 marks) 8. Using molecular orbital theory, write the electronic configurations of O2, O2+ and O2-, and calculate their bond orders. Arrange them in order of increasing bond length and state which are paramagnetic. (3 marks) 9. Predict the shape and the hybridisation of the central atom in each: BeCl2, BF3, NH3, PCl5, SF6, H2O. Give the number of bond pairs and lone pairs for each. (3 marks) 10. The dipole moment of NH3 is 1.47 D and that of NF3 is 0.24 D, although both are pyramidal. Explain the difference using the direction of bond moments and the lone pair moment. Why is the dipole moment of BF3 zero? (3 marks) Section C - Diagrams (10 marks) 11. Draw the shapes of the sp, sp2 and sp3 hybrid orbitals and show their orientations with bond angles. (4 marks) 12. Draw the molecular orbital energy level diagram of O2 and use it to show why O2 is paramagnetic. (3 marks) 13. Draw the Lewis dot structures of CO3 2- and its resonance structures. (3 marks) Section D - Analysis and Higher-order Thinking (15 marks) 14. Explain how VSEPR theory predicts the shapes of molecules. Why is the bond angle in H2O smaller than in NH3, which is smaller than in CH4? (5 marks) 15. Explain the concept of hybridisation with examples of PCl5 and SF6. Why are the axial bonds in PCl5 longer than the equatorial bonds? (5 marks) 16. Compare valence bond theory and molecular orbital theory. Using MOT, explain why He2 does not exist and why N2 has a higher bond enthalpy than O2. (5 marks) Instructions: Time allowed 2 hours. Attempt all sections. Draw diagrams neatly. Show the steps of all calculations with units. Write formulae with plain digits (e.g. SF6).