1. Chemical Bonding Forming compounds from atoms

2. Intramolecular Interactions Intramolecular = inside the molecules. – The bonds that form between the atoms. Metallic Ionic Covalent – Non polar and polar

3. Octet rule Atoms form bonds in order to have a full outer shell. i.e. they have the same electron configuration as the noble gases One atom can give electrons away (cation) and another formally receive electrons (anion) to form Ionic bonds Or they can share the electrons to form Covalent bonds Or the metals can live in a ‘sea of electrons’ where the electrons are free to move about. This give rise to Metallic bonding

4. WARNING I know you’re clever, however this is a beginner chemistry paper. There will be exceptions to the broad rules that will be used here as you move on in chemistry. It is too time consuming at this level to cover each and every exception to the general rules. (I’m happy to talk to you about them, but the exam will be based on the generic, not the exceptions)

5. Bonding The periodic table can be split into three parts based on electronegativity (how much they want another electron). High values, Medium values and Low values. The difference in these values dictate what kind of bonding will be shared between two atoms.

6. Bonding High electronegativity

7. Bonding Medium electronegativity

8. Bonding Low electronegativity

9. Bonding When two metals in the blue section bond, Metallic Bonding is the result

10. Examples: Na (s) Pd (s) Ni (s) etc etc…

11. Bonding When an element from the blue and one from the red area bond, Ionic bonding results

12. Examples: Na and Cl Zn and F Ni and O etc etc…

13. Covalent Bonding The next two discussed are when the electrons are shared in a two electron bond The electrons can either be shared evenly (non-polar covalent) where the two atoms have similar electronegativity or unevenly (polar covalent) where one atom is more electronegative than the other

14. Bonding When two elements from the orange area bond, Non-Polar Covalent results

15. Examples: C and H C and P S and H etc etc…

16. Bonding When one element from the orange area bonds with one from the red, a Polar Covalent results

17. Electronegativity H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.1 O 3.5 F 4.0 Na 1.0 Mg 1.3 Al 1.5 Si 1.8 P 2.1 S 2.4 Cl 3.0 K 0.9 Ca 1.1 Ga 1.8 Ge 2.0 As 2.2 Se 2.5 Br 2.8 Rb 0.9 Sr 1.0 In 1.5 Sn 1.7 Sb 1.8 Te 2.0 I 2.2 Cs 0.9 Ba 0.9 Tl 1.5 Pb 1.6 Bi 1.7 Po 1.8 At 2.0 Fr 0.9 Ra 0.9 The reason these bonds are polar is that their electronegativity differs by > 0.5 e.g. Si – Br 2.8 – 1.8 = 1.0 therefore polar

18. Examples: Cl and H O and P S and F C and O etc etc…

19. What are these? Metallic, Ionic, Covalent (polar or non-polar) 1.Mg and Ca 2.K and Cl 3.C and H 4.C and O 5.As and Br 6.O and O 7.Ti and Cr 8.Na and F 9.Si and I 10.Sb and Cl

20. What are these? Metallic, Ionic, Covalent (polar or non-polar) 1.Mg and CaMetallic 2.K and ClIonic 3.C and HCovalent non-polar 4.C and OCovalent polar 5.As and BrCovalent polar 6.O and OCovalent non-polar 7.Ti and CrMetallic 8.Na and FIonic 9.Si and ICovalent polar 10.Sb and ClCovalent polar

21. Lewis dot diagrams To work out the bonding in a covalent molecule, Lewis dot diagrams can be used Most elements obey the octet rule, where they want 4 pairs of electrons By following some basic rules, structures can be worked out

22. Instructions: 1.Draw each atom with its valence electrons represented by dots around the symbol. 2.Underneath the symbol note how many electrons each atom needs to share to achieve the stable octet. (Remember hydrogen is an exception to the rule). 3.Work out how to fit the atoms together so each has the appropriate number of electrons. A good rule of thumb is that the atom that needs to share the most electrons is probably the central atom (the atom the others are bonded to). 4.Draw a bond diagram of the molecule including lone pairs of electrons on the central atom.

23. Example: Cl 2 1.Draw each atom with its valence electrons represented by dots around the symbol. Chlorine had 7 valence electrons from the periodic table. Therefore 3 pairs and one remaining

24. Example: Cl 2 2.Underneath the symbol note how many electrons each atom needs to share to achieve the stable octet. Needs on more electron 3.Work out how to fit the atoms together so each has the appropriate number of electrons. A good rule of thumb is that the atom that needs to share the most electrons is probably the central atom (the atom the others are bonded to). These could pair

25. Example: Cl 2 4.Draw a bond diagram of the molecule including lone pairs of electrons on the central atom. This has formed a single, 2-electron bond: Cl - Cl

26. Example: CH 4 1.Draw each atom with its valence electrons represented by dots around the symbol. Carbon had 4 valence electrons from the periodic table. Therefore 4 individual electrons and H has one

27. Example: CH 4 2.Underneath the symbol note how many electrons each atom needs to share to achieve the stable octet. C needs 4 x1 more electrons H needs 1 x4 more electrons

28. Example: CH 4 3.Work out how to fit the atoms together so each has the appropriate number of electrons. A good rule of thumb is that the atom that needs to share the most electrons is probably the central atom (the atom the others are bonded to). These could pair

29. Example: CH 4 4.Draw a bond diagram of the molecule including lone pairs of electrons on the central atom. This has formed a four single, 2-electron bonds:

30. More examples: CO 2 CH 2 O SiF 4 O 2 NH 3

31. VSEPR Valence shell electron pair repulsion. – Gives the shape of the molecule that can have important consequences for some properties (Intermolecular interactions) – Have to work out how many electron density areas (‘steric number’) and lone pairs there are using Lewis dot diagrams

33. Molecular Polarity Polar bonds can lead to molecules being polar overall. Must be careful that the polarities do not cancel you by geometry – e.g. Label as polar or non-polar by determining geometry: CO 2 CCl 4 H 2 O PCl 3

34. Molecular Polarity Polar bonds can lead to molecules being polar overall. Must be careful that the polarities do not cancel you by geometry – e.g. Label as polar or non-polar by determining geometry: CO 2 Linear, therefore cancel out and non-polar CCl 4 Tetrahedral, therefore cancel out and non-polar H 2 O Bent, not cancelling therefore polar PCl 3 Trigonal pyramidal, not cancelling therefore polar

35. More examples CF 2 O CS 2 NCl 3 HBr CH 2 CH 2 SCl 2 CH 2 F 2