1. Topic 5: The Chemical Bonding Chemical bonding is a force that holds atoms together to give molecules Ionic bonding  Ionic bonds are generally formed when a metal from the left side of the periodic table (low ionisation energy) gives its electrons to a non-metal from the upper right part of the periodic table (high electron affinity).  Complete transfer of electrons from one atom to another to form a positive and negatively charged ions

2. Na(s) + 1/2 Cl 2 (g)  Na + + Cl -  NaCl (s) NON-DIRECTIONAL bonding via Coulomb (charge) interaction

3. Properties of ionic bonds  Ionic crystal lattice is built up of ions that have opposite charges.  Electrostatic forces of attraction extend in all directions. A huge variety of positive and negative ions alternate with each other, and giant lattices are formed. Ionic bonds are nondirectional.

4.  Ionic compounds are hard, and have high melting points and boiling points. The strong electrostatic force between the ions is the reason for these characteristics.  Ionic compounds are soluble in polar solvents, and they are insoluble in nonpolar solvents.  Ionic compounds are conductors in the fused state and solution. In the solid state, these compounds are poor conductors of electricity.  Reactions of ionic compounds are reactions of their ions, and they occur instantly in solution

5. Covalent bond and molecules Between nonmetallic elements of similar electronegativity. Formed by sharing electron pairs Stable non-ionizing particles, they are not conductors at any state Examples; O 2, CO 2, C 2 H 6, H 2 O, SiC

6. Bond types in covalent molecules

7. Pi (π) – bond- side to side overlap of p-orbitals Two p-orbitals of carbon

8. 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. In ammonia, N is central Building a Dot/ Lewis Structure Ammonia, NH 3

9. 4. If any electron left, put them at the central atom as lone pair 5. If the central atom is not completely filled, form a double bond using lone pair electrons of the peripheral atom (see example below) 3. Distribute starting the peripheral atoms to completely fill them.

10. 6. Decide the charge of the atoms  Check the electrons on each atom. If more than the valance electron, put negative charge if less put positive charge otherwise neutral

11. Lewis formula of methane- C= 4 valance electrons H= 4 electrons (one for each hydrogen) Total= 8 electrons to be distributed

12. Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 6 + 18 + 2 = 26 e- Step 1. Central atom = S 10 pairs of electrons are left. Sulfite ion, SO 3 2- Step 3. Form sigma bonds (skeleton)

13. Remaining pairs become lone pairs, first on outside atoms then on central atom. Sulfite ion, SO 3 2- Each atom is surrounded by an octet of electrons. OO O S NOTE - must add formal charges (O -, S + ) for complete dot diagram Make one double bond between oxygen and sulfur

15. Dative bonding Dative (co-ordinate) bonding is a type of covalent bond in which one of the atoms provides both electrons to form the covalent bond. Example

16. Other examples: In your reading material

18. Molecular Geometry Molecular geometry is the three dimensional arrangement of atoms in a molecule. There is a simple method called the Valence Shell Electron Pair Repulsion theory (VSEPR for short)

19. VSEPR Rules Rule 1 Lone pair of electrons repels neighbouring lone pair more than bonding electron pairs. Order of repulsion: Lone pair-lone pair > Lone pair-Bonding pair > Bonding pair-bonding pair Strongest weakest

20. Rule 2 Double and triple bonds in a molecule are treated like single bonds. Rule 3 If a molecule has two or more resonance structures, we can apply the VSEPR theory to any one of them.

21. MoleculeLewis StructureNumber of electron pairs CH 4 NH 3 SHAPE Tetrahedral Trigonal Pyramidal 4 4 (3 shared 1 lone pair)

22. MoleculeLewis StructureNumber of electron pairs H2OH2O CO 2 SHAPE Bent or V 4 (2 shared 2 lone pairs) 2 Linear

23. MoleculeLewis StructureNumber of electron pairs BeCl 2 BF 3 SHAPE 2 3 Linear Trigonal Planar

24. Hybridization of atomic orbitals Hybridization is the mixing of pure atomic orbitals of similar energy to give equal number of hybrid orbitals of equal energy and similar properties.

25. Electron configuration of carbon Has two unpaired electrons Means it can only overlap with two H atoms to form CH 2 (it is unstable!) In reality it forms CH 4 It means there four unpaired electrons of carbon which overlap with four H atoms to form CH 4  One electron from 2S jumps to 2pz

26. Now, the four unpaired electron overlaps with 4 H atoms to form 4 C-H bonds But, experiment shows the 4 C-H bonds are identical. That implies the bonds are formed from overlap of identical orbitals that is where the concept of hybridization of the orbitals of carbon comes

27. The 2s orbital and the three p-orbitals mix to form four identical orbitals called SP 3 (to say that one s and 3 p orbitals are mixed) SP 3 hybridized orbital has 25% s-character and 75% p-character Each of the four SP 3 orbitals contain one electron which then overlap with the four hydrogen to form CH 4 These bonds are formed by the overlap of SP 3 hybrid orbital and s-orbital

28. Pi bond by side to side overlap of the p-orbitals The remaining 2 p-orbitals and one 2s orbital mixt to give a hybridized orbital called SP 2 (this orbital has 33.3% s-character and 66.6% p-character)

29. Acetylene Pi-bond 2-pi bonds which are formed from side to side overlap between p-orbitals of the carbon atoms One s and one p-orbitals mix to give an orbital called sp hybridized orbital (50% s-character and 50% p-character) Each carbon has two sp-orbitals which they overlap with each other to form the sigma bond between them and the other sp orbital overlap with 1s orbital of hydrogen to form the sigma bonds between carbon and hydrogen

30. Intermolecular forces  between molecules and strongly influence the physical properties such as the boiling point, viscosity etc. i)Londons/ dispersion force (weak force)- due to instantaneous dipole-induced dipole the ease with which the electron distribution in the atom or molecule can be distorted-determines the property

31. ii) Dipole–dipole forces Interaction between two dipolar molecules iii) Ion-Dipole Forces NaCl in water S O O +ve-pole -ve pole S O O +ve-pole -ve pole Attract each other

32. When hydrogen is involved it is called hydrogen bonding The Hydrogen Bond- STRONGEST INTERMOLECULAR FORCE Water in the liquid and solid states exists as groups in which the water molecules are linked together by hydrogen bonds.

34. Questions taken from TUTORIAL 3 Explain why? i) Ethanol (CH 3 CH 2 OH), boils at 78 o C and dimethyl ether (CH 3 -O- CH 3 ) boils at –24 O C. ii) Methanol (CH 3 OH) is miscible (mixes) with water whereas benzene is not iii) Solid NaCl does not conduct electricity but an aqueous solution of NaCl does. iv) H2O boils at 100 o C at room pressure where as NH 3 boils at – 33 o C.