The Chemical Bond

1. The Chemical Bond

2. What is a chemical bond? • It is a force of attraction between atoms, ions or molecules.

3. How are bonds formed ? • By electron loss, gain or sharing.

4. Which elements lose electrons? • Metals • Cations are formed

5. Why metals lose electrons ? • They usually have one, two or three electrons. • By losing these electrons they can achieve a full outer shell and become more stable. • It is in energy terms easier to lose electrons than gain electrons

6. Which elements gain electrons? • Non-metals • Anions are formed

7. Why non-metals gain electrons? • They usually have five, six or seven electrons. • By gaining three, two or one electron, they can achieve a full outer shell and become more stable. • In energy terms it is easier to gain these electrons than lose the outer electrons.

8. Bonds • Covalent • Polar covalent • Dative covalent • Giant covalent • Ionic • Metallic

9. How are covalent bonds formed? By sharing electrons between non-metals.

10. Dot and cross diagram-Covalent bond • Each H atom contributes one electron to the shared pair. • Each H atom has two electrons in its outer shell. • Each H atom now has a full outer shell H : H

11. What is electronegativity? • It is the ability of an atom to attract the bonding electrons in a bond.

12. Electronegativity values • Hydrogen 2.1 Lithium 1.0 • Beryllium 1.5 Boron 2.0 • Carbon 2.5 Nitrogen 3.0 • Oxygen 3.5 Fluorine 4.0

13. Which element is most electronegative? • Fluorine • The electronegativity is 4.0

14. How does the presence of a more electronegative atom affect a covalent bond? • The more electronegative atom attracts the shared electrons more towards itself. • The more electronegative atom acquires a partial negative charge. • The less electronegative atom a partial positive charge. • The bond becomes a polar covalent bond.

15. Calculate the electronegativity difference for the following • C – H . Li - F • O – H • H – F • N – H • C –N • C – O • C - C

16. Which bond has the highest electronegativity difference? • Li F • Electronegativity differences greater than or equal 1.6 show that the compound is ionic • This compound is ionic!

17. Ionic or polar covalent ? • This depends on the electronegativity difference • When the difference is relatively small, the bonding is polar covalent. • When the difference is large, the bonding becomes ionic. (1.6 or more)

18. How are ionic bonds formed ? • By exchanging electrons. • Metals lose electrons and become positive ions. • Non-metals gain electrons and become negative ions. • The oppositely charged ions attract each other. This is ionic bond.

19. Ionic bonding

20. Ionic bonding

21. Giant Covalent Bond • Structure of diamond

22. Giant Covalent Bond • Structure of diamond

23. Giant Covalent Bond • Structure of Graphite

24. How are metallic bonds formed? • In metals , the outer electrons are free. • These delocalised electrons form a ‘sea’ of electrons. • The metal atoms exist as positive ions immersed in this ‘sea’ of electrons. • The attraction between the positive ions and the ‘sea’ of electrons is known as the metallic bond.

25. Metallic bonding • Metallic bonding in potassium – the outermost electrons are ‘free’ and form a ‘sea’ of electrons which attracts the K+ions.

26. Intermolecular Bonds . Van der Waals forces . Dipole dipole attraction . Hydrogen bonds

27. Van der Waals Forces • Random electron movements creates dipoles. • Dipoles induces dipoles. • Process is repeated. • Oppositely charged dipoles attract.

28. Van der Waals Forces

30. Van der Waals Forces- Factors • More electrons, higher forces • Larger the area of contact, higher forces • Linear molecules, higher forces • Branched chains, lower forces • Spherical molecules, lower forces

31. Dipole-dipole attraction . More electronegative atom in a covalent bond. . Permanent dipoles. . Attraction between dipoles.

32. Dipole-dipole attraction-Factors • More the number of electronegative atoms, higher attraction • More electronegative atoms, higher attraction

33. Hydrogen Bond . Hydrogen in a covalent bond with nitrogen, oxygen or fluorine. . Permanent dipoles. . Attraction between dipoles. . Strongest intermolecular force. . Intermolecular forces much weaker than normal covalent bonds

34. Hydrogen bonding in water

35. H-bond in H2O • But for its H-bonding water would be a gas at room temperature. • Due to H-bonding ice floats on water. • Water freezes with a lot of space in between due to H-bonding. This makes ice lighter. H O||||||||||| H H H O O|||||||| H H

36. Hydrogen bonding in water

37. Hydrogen bonding ….. • Ammonia can form H-bonds. • HF can also form H-bonds.

38. Trend in the melting and boiling point of Group IV hydrides mp • As the elements get larger, there are more electrons. • Van der Waals force is larger. • MP and BP is larger from CH4 to PbH4. CH4 PbH4

39. Anomalous melting and boiling points of hydrides • Why ammonia, the lightest hydride has the highest melting and boiling point than the corresponding hydrides in Group V? • Why water, the lightest hydride in GroupVI has the highest melting and boiling points? • Why hydrogen fluoride, the lightest hydride in Group VII has the highest melting and boiling points?

40. What are these hydrides • Gr V: NH3, PH3, AsH3, SbH3,BiH3 • Gr VI: H2O, H2S, H2Se, H2Te, H2 Po • Gr VII: HF, HCl, HBr, HI, HAt

41. Why the mp or bp of NH3 is unusually so high? • Nitrogen is electronegative. • NH bonds are polar • H-bonds formed between ammonia molecules. • No H-bonds in other molecules, only Van der Waals Forces.Partially covalent ionic bonds . High charge density of positive ions. . Positive ions polarising negative ions. . Resulting electron sharing – partial covalency.Sigma and pi-bonds . Sigma bond – direct overlap of orbitals . Pi- bond – sideways overlap of orbitals. N-H||||||| N -H |||| H-bond BP BiH3 NH3

42. What are dative covalent bonds? • In a dative covalent bond, both the shared electrons come from the same atom. • The atom giving the electron pair is the donor. • The atom receiving the electron pair is the acceptor. • There is no difference in length and strength between a normal and a dative covalent bond.

43. The End