1. Chapter 8 “Covalent Bonding” Part 3 Ball-and-stick model

2. VSEPR Theory: stands for... Valence Shell Electron Pair Repulsion Predicts the three dimensional shape of molecules. The name tells you the theory: –Valence shell = outside electrons. –Electron Pair repulsion = electron pairs try to get as far away as possible from each other. Can determine the angles of bonds.

3. VSEPR Based on the number of pairs of valence electrons, both bonded and unbonded. Unbonded pair also called lone pair. CH 4 - draw the structural formula Has 4 + 4(1) = 8 wants 8 + 4(2) = 16 (16-8)/2 = 4 bonds

4. VSEPR for methane (a gas): Single bonds fill all atoms. There are 4 pairs of electrons pushing away. The furthest they can get away is 109.5º CHH H H This 2-dimensional drawing does not show a true representation of the chemical arrangement.

5. 4 atoms bonded Basic shape is tetrahedral. A pyramid with a triangular base. Same shape for everything with 4 pairs. C HH H H 109.5º

6. Ammonia (NH 3 ) = 107 o Water (H 2 O) = 105 o Carbon dioxide (CO 2 ) = 180 o Refer to Molecular Geometry Chart

7. Methane has an angle of 109.5 o, called tetrahedral Ammonia has an angle of 107 o, called pyramidal Note the unshared pair that is repulsion for other electrons.

8. Hybrid Orbitals The VSEPR theory works well when accounting for molecular shapes, but it does not help much in describing the types of bonds formed. Orbital hybridization provides information about both molecular bonding and molecular shape.

9. A bond can result from head to head overlap of atomic orbitals on neighboring atoms.

10. Hybrid Orbitals In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals SP 3 Hybridization : four pairs of bonding electrons. 1 s and 3 p orbitals SP 2 Hybridization : three pairs of bonding electrons. 1 s and 2 p orbitals SP Hybridization: two pairs of bonding electrons 1 s and 1 p orbtial.

11. Multiple Bonds Everything we have talked about so far has only dealt with what we call sigma bonds Sigma bondA bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms. Sigma bond (  )  A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.

12. Pi bondA bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis). Pi bond (  )  A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).

13. Example: H 2 C=CH 2

15. Example: HC  CH

16. Bond Polarity Consider HCl H = electronegativity of 2.1 ; Cl = electronegativity of 3.0 the bond is polar the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge A molecule that has two poles is called a dipolar molecule, or dipole.

17. Molecular Polarity Molecular Polarity depends on: 1. The relative electronegativities of the atoms in the molecule – bond plarity 2. The shape of the molecule - Molecules that have symmetrical charge distributions are usually non- polar If there are equal polar bonds that balance each other around the central atom, then the overall molecule will be NONPOLAR with no dipole moment, even though the bonds within the molecule may be polar.

18. Molecular Polarity - Polar bonds cancel - There is no dipole moment - Molecule is non-polar - Polar bonds do not cancel - There is a net dipole moment - The molecule is polar

19. Polar molecules carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

20. Polar molecules water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e - away from H = very polar!

21. Intermolecular Forces Intermolecular forces are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature.

22. Intermolecular Forces The weakest attractions is van der Waals forces and there are two types : 1. Dispersion forces : weakest of all, caused by motion of e - ; increases as # e - increases halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A 2. Dipole Interactions : Occurs when polar molecules are attracted to each other. Slightly stronger than dispersion forces. Opposites attract, but not completely hooked like in ionic solids.

23. #2. Dipole Interactions    

24. Intermolecular Forces Hydrogen Bonds : It is the attractive force caused by hydrogen bonded to N, O, F N, O, and F are very electronegative, so this is a very strong dipole. Hydrogen bonds are extremely important in determining the properties of water and biological molecules such as proteins. This is the strongest of the intermolecular forces. Hydrogen bond

25. Hydrogen bonding allows H 2 O to be a liquid at room conditions. H H O H H O H H O H H O H H O H H O H H O

26. Order of Intermolecular attraction strengths 1)Dispersion forces are the weakest 2)A little stronger are the dipole interactions 3)The strongest is the hydrogen bonding 4)All of these are weaker than ionic bonds