Covalent Bonding

1. 8 Covalent Bonding 8.1 Formation of Covalent Bonds 8.2 Dative Covalent Bonds 8.3 Bond Enthalpies 8.4 Estimation of Average Bond Enthalpies using Data from Energetics

2. 8.5 Use of Average Bond Enthalpies to Estimate the Enthalpy Changes of Reactions 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii 8.7 Shapes of Covalent Molecules and Polyatomic Ions 8.8 Multiple Bonds 8.9 Covalent Crystals

3. 8.1 Formation of Covalent Bonds

4. Attraction between oppositely charged nuclei and shared electrons ( _____________ in nature) H H Shared electrons e- e- 8.1 Formation of Covalent Bonds (SB p.213) A. Electron Sharing in Covalent Bonds electrostatic The shared electron pair spends most of the time between the two nuclei. Overlapping of atomic orbitals  covalent bond formation

5. 8.1 Formation of Covalent Bonds (SB p.213) A hydrogen molecule is achieved by partial overlapping of 1s orbitals

6. 8.1 Formation of Covalent Bonds (SB p.214) Electron density map for covalent compounds There is substantial electron density at all points along the internuclear axis. Compare electron-density-map for ionic compounds: Thus electrons are shared between the two atoms.

7. 8.1 Formation of Covalent Bonds (SB p.214) B. Covalent Bonds in Elements • Hydrogen molecule Dot and cross diagram

8. 8.1 Formation of Covalent Bonds (SB p.215) • Chlorine molecule • Oxygen molecule

9. 8.1 Formation of Covalent Bonds (SB p.215) • Nitrogen molecule

10. 8.1 Formation of Covalent Bonds (SB p.216) C. Covalent Bonds in Compounds

11. 8.1 Formation of Covalent Bonds (SB p.216)

12. 8.1 Formation of Covalent Bonds (SB p.216 – 217) D. Octet Rule and its Limitations In forming chemical bonds, atoms tend to achieve the stable noble gas electronic configuration with 8 electrons in the valence shell (except helium which has 2 electrons in the valence shell) by gaining, losing or sharing of electrons.

13. not fullfilling octect (electron deficient) electrons from F 8.1 Formation of Covalent Bonds (SB p.217) 1. Boron Trifluoride (BF3) B: small atomic size high I.E.’s required to become a cation. Why doesn’t B form ionic compounds with F?

14. 8.1 Formation of Covalent Bonds (SB p.207) electrons from Cl 2. Phosphorus Pentachloride (PCl5) There is low-lyingvacant d-orbital in P. Check Point 8-1 Why Phosphorus can expand its octet to form PCl5?

15. 8.2 Dative Covalent Bonds

16. 8.2 Dative Covalent Bonds (SB p.218) Dative Covalent Bonds A dative covalent bond is formed by the overlapping of an empty orbital of an atom with an orbital occupied by a lone pair of electrons of another atom. Remarks(1) The atom that supplies the shared pair of electrons is known as the donor while the other atom involved in the dative covalent bond is known as the acceptor. (2) Once formed, a dative covalent bond cannot bedistinguished from a ‘normal’ covalent bond.

17. 8.2 Dative Covalent Bonds (SB p.218 – 219) A. NH3BF3 molecule

18. 8.2 Dative Covalent Bonds (SB p.219) B. Ammonium Ion (NH4+)

19. 8.2 Dative Covalent Bonds (SB p.219 – 220) D. Aluminium Chloride Dimer (Al2Cl6) Al: relative small atomic size; high I.E.’s required to become a cation of +3 charge. AlCl3

20. 8.2 Dative Covalent Bonds (SB p.219 – 220) D. Aluminium Chloride Dimer (Al2Cl6) (a dimer of AlCl3) Why doesn’t Al form ionic compounds with Cl? Check Point 8-2

21. 8.3 Bond Enthalpies

22. 8.3 Bond Enthalpies (SB p.220) Bond Enthalpy Bond enthalpy is the energy associated with a chemical bond. When a chemical bond is broken or formed, a certain amount of energy is absorbed from or released to the surroundings.

23. 8.3 Bond Enthalpies (SB p.220) • Example: • Combustion of methane

24. 8.3 Bond Enthalpies (SB p.221) Standard enthalpy changes of combustion of the homologous series of alkanes and alkanols

25. e.g. H-H(g)  2H(g) H = +431 kJ mol-1 ø 8.3 Bond Enthalpies (SB p.221) Bond Dissoication Enthalpy Bond dissociation enthalpy is the enthalpy change when one mole of a particular bond in a particular environment is broken under standard conditions.

26. ø CH4(g) CH3(g) + H(g) H= +422 kJ mol-1 CH3(g) CH2(g) + H(g) H = +480 kJ mol-1 CH2(g) CH(g) + H(g) H= +425 kJ mol-1 CH(g) C(g) + H(g) H = +335 kJ mol-1 ø ø ø 8.3 Bond Enthalpies (SB p.221) Why do successive B.D.E. of C-H differ? (Average) bond enthalpy; E(C-H) = +415.5 kJ mol-1

27. 8.3 Bond Enthalpies (SB p.222) Average Bond Enthalpies Average bond enthalpy is the average of the bond dissociation enthalpies required to break a particular chemical bond. Let's Think 1

28. 8.4 Estimation of Average Bond Enthalpies using Data from Enegetics

29. 8.4 Estimation of Average Bond Enthalpies using Data from Enegetics (SB p.223) A. Derived from the Enthalpy Change of Atomization of a Compound Atomization of a compound means the breaking down of one mole of the gaseous compound into its constituent atoms in the gaseous state.

30. ø C(g) + 4H(g) H= +1 662 kJ mol-1 8.4 Estimation of Average Bond Enthalpies using Data from Enegetics (SB p.223) • Example: • Atomization of methane The atomization of methane involves the breaking of a four C-H bonds. Assume that all four C-H bonds are equal in strength. The average bond enthalpy of C-H bonds = ¼ x (+1 662) kJ mol-1 = +415.5 kJ mol-1 E(C-H) = +415.5 kJ mol-1

31. 8.4 Estimation of Average Bond Enthalpies using Data from Enegetics (SB p.223) • Two ways to determine the enthalpy change of atomization of methane: • 1. From successive bond dissociation enthalpies • 2. From enthalpy cycle and Hess’s law

32. 8.4 Estimation of Average Bond Enthalpies using Data from Enegetics (SB p.224 – 225) B. Derived from the Enthalpy Changes of Atomization of Two Compounds The enthalpy change of atomization of butane (C4H10) and pentane (C5H12) are +5165 kJ mol-1 and +6337 kJ mol-1 respectively. Find a values for the bond enthalpies of C-H and C-C based on the above data.

33. 8.4 Estimation of Average Bond Enthalpies using Data from Enegetics (SB p.224 – 225) B. Derived from the Enthalpy Changes of Atomization of Two Compounds For butane, 3 E(C-C) + 10 E(C-H) = +5 165 kJ mol-1 ….(1) For pentane, 4 E(C-C) + 12 E(C-H) = +6 337 kJ mol-1 ….(2) Solving simultaneous equations (1) and (2), we obtain the following bond enthalpy values. E (C-H) = +412.25 kJ mol-1 E (C-C) = +347.5 kJ mol-1

34. 8.5 Use of Average Bond Enthalpies to Estimate the Enthalpy Changes of Reactions

35. Sum of bond enthalpies of reactants Enthalpy change of reaction - = 8.5 Use of Average Bond Enthalpies to Estimate the Enthalpy Changes of Reactions (SB p.225) Reaction between ethene and hydrogen Sum of bond enthalpies of products

36. 8.5 Use of Average Bond Enthalpies to Estimate the Enthalpy Changes of Reactions (SB p.226) Enthalpy level diagram for the reaction between ethene and hydrogen

37. 8.5 Use of Average Bond Enthalpies to Estimate the Enthalpy Changes of Reactions (SB p.226) Reaction between methane and oxygen Check Point 8-5

38. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii

39. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.227) A. Bond Enthalpies as an Indication of the Strength of Covalent Bonds • Gives a direct measure of the strength of a covalent bond It is the energy required to break the bond • Not in proportion to the bond order(The number of bonding electrons divided by two)

40. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.228) B. Bond Lengths • The distance between the two bonded nuclei • Inversely related to bond strength • Not constant • Depends on the local environment of that particular bond • Determined experimentally by electron diffraction, X-ray diffraction or spectroscopic techniques

41. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.228) C. Relationship between Bond Lengths and Bond Enthalpies Any conclusion for the relationship between bond length & bond enthalpy? Usually a longer bond length corresponds to a lower value of bond enthalpy (weaker bond).

42. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.228) Special Situation for F2 BondBond Length /nmBond Enthalpy / kJ mol-1 F-F 0.142 158 Cl-Cl 0.199 242 Br-Br 0.228 193 I-I 0.266 151 Explain why the bond enthalpy of F-F is smaller than that of Cl-Cl even though the bond length of F-F is the shortest among the halogens.

43. Non-bonding e-/ lone pair of e- F F 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.228) As the size of fluorine atom is very small, the repulsion between the non-bonding pairs of electrons on the fluorine atoms weaken the F-F bond.

44. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.228) D. Covalent Radii • Half the internuclear distance between two atoms in a covalently bonded molecule • Generally taken as half of the bond length of homoatomic covalent molecules (where identical atoms are bonded together)

45. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.228) The covalent radius of an atom is taken as half of the bond length of a homoatomic molecule

46. 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.229) The covalent radii (in nm) of some elements

47. Bond length of a covalent bond A-B Covalent radius of atom A Covalent radius of atom B = + 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.229) Predicting bond length of A-B if rA & rB are known

48. By what technique can the bond lengths be determined experimentally? Similarelectronegativity 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.229 – 230) Calculated and experimentally determined bond length

49. Quite differentelectronegativity 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii (SB p.229 – 230) Calculated and experimentally determined bond length Check Point 8-6 Let's Think 2

50. 8.7 Shapes of Covalent Molecules and Polyatomic Ions