2. Electronegativity ++ –– 00 00 HClHH

3. Electronegativity Electronegativity describes how electrons are shared in a compound Consider the compound HCl The electron clouds represent where the two electrons in the HCl bond spend their time (sizes of atoms are not being shown) The shared electrons spend more time around Cl than H. In other words Cl is more electronegative than H. HCl ++ ––

4. Electronegativity table These numbers are derived from several factors including EA, IE, atomic radius You do not need to understand where the numbers come from You need to know that a high number means the element has a greater pull on electrons You will also need to calculate the difference between values to estimate the % ionic or % covalent character of a bond

5. Electronegativity vs Polarity of bond

6. Calculating EN differences The first step in defining the polarity of a bond is to calculate electronegativity difference (  EN)  EN = EN large - EN small E.g. for NaCl,  EN = 2.9 - 1.0 = 1.9 Next, estimate from fig 7.12 the % ionic character: about 65% (60 - 70%) Q-Give the % ionic character for MgO, CH, HCl MgO = 3.5 - 1.3 = 2.2 … 80% ionic (75-85) CH = 2.5 - 2.1 = 0.4 … 7% ionic (5 - 10) HCl = 2.9 - 2.1 = 0.8 … 20 % ionic (15-25) Note if % ionic is 20%, then % covalent is 80%

7. Defining polarity For our purposes we will define polarity in the following fashion: 0-10 % is non-polar, 10- 50% is polar (covalent), 50%+ is ionic This is a crude estimate. In reality, the only non-polar bond between 2 atoms occurs in diatomic molecules (O 2 :  EN = 3.5 - 3.5 = 0) Q - what is the polarity of the bonds in MgO, CH, HCl? MgO = 80% ionic = ionic CH = 7% ionic = non-polar HCl = 20 % ionic = polar covalent

8. Polarity in molecules A bond that you calculate to be polar may not be polar if the molecule is symmetrical Imagine a tug-of-war between atoms of the same strength around a central atom The pull in one direction is the bond polarity or “bond dipole”. The overall/molecular polarity is also known as “dipole moment” If the pull is the same from all directions then the electrons are not attracted to one atom over another and the molecule is non-polar

9. Polarity in molecules: Examples NH 3 NH 4 + BrCl Lewis Structure  EN of bonds Bond polarity Symmetrical molecule? Polarity of the molecule

10. Why do some solids dissolve in water but others do not? Why are some substances gases at room temperature, but others are liquid or solid? What gives metals the ability to conduct electricity, what makes non-metals brittle? The answers have to do with … Intermolecular forces Questions

11. Overview There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces We have already looked at intramolecular bonds (ionic, polar, non-polar) Intermolecular forces (IMF) have to do with the attraction between molecules (vs. the attraction between atoms in a molecule) IMFs come in six flavours: 1) ionic, 2) dipole - dipole, 3) H-bonding, 4) London forces, 5) covalent (network solids), 6) metallic Intermolecular forces

12. Ionic, Dipole - Dipole attractions We have seen that molecules can have a separation of charge This happens in both ionic and polar bonds (the greater the EN, the greater the dipoles) HCl ++ –– Molecules are attracted to each other in a compound by these +ve and -ve forces ++ –– ++ –– ++ –– ++ ––

13. H - bonding H-bonding is a special type of dipole - dipole attraction that is very strong It occurs when N, O, or F are bonded to H Q- Calculate the  EN for HCl and H 2 O A- HCl = 2.9-2.1 = 0.8, H 2 O = 3.5-2.1 = 1.4 The high  EN of NH, OH, and HF bonds cause these to be strong forces (about 5x stronger than normal dipole-dipole forces) They are given a special name (H-bonding) because compounds containing these bonds are important in biological systems

14. London forces Non-polar molecules do not have dipoles like polar molecules. How, then, can non-polar compounds form solids or liquids? London forces are named after Fritz London (also called van der Waal forces) London forces are due to small dipoles that exist in non-polar molecules Because electrons are moving around in atoms there will be instants when the charge around an atom is not symmetrical The resulting tiny dipoles cause attractions between atoms/molecules

15. London forces Instantaneous dipole:Induced dipole: Eventually electrons are situated so that tiny dipoles form A dipole forms in one atom or molecule, inducing a dipole in the other

16. Testing concepts 1.Which attractions are stronger: intermolecular or intramolecular? 2.How many times stronger is a covalent bond compared to a dipole-dipole attraction? 3.What evidence is there that nonpolar molecules attract each other? 4.Which chemical has the weakest intermolecular forces? Which has the strongest? How can you tell? 5.Suggest some ways that the dipoles in London forces are different from the dipoles in dipole-dipole attractions. 6.A) Which would have a lower boiling point: O 2 or F 2 ? Explain. B) Which would have a lower boiling point: NO or O 2 ? Explain.

17. 7.Which would you expect to have the higher melting point (or boiling point): C 8 H 18 or C 4 H 10 ? Explain. 8.What two factors causes hydrogen bonds to be so much stronger than typical dipole-dipole bonds? 9.So far we have discussed 4 kinds of intermolecular forces: ionic, dipole-dipole, hydrogen bonding, and London forces. What kind(s) of intermolecular forces are present in the following substances: a) NH 3, b) SF 6, c) PCl 3, d) LiCl, e) HBr, f) CO 2 (hint: consider  EN and molecular shape/polarity) Challenge: Ethanol (CH 3 CH 2 OH) and dimethyl ether (CH 3 OCH 3 ) have the same formula (C 2 H 6 O). Ethanol boils at 78  C, whereas dimethyl ether boils at -24  C. Explain why the boiling point of the ether is so much lower than the boiling point of ethanol..

18. Testing concepts 1.Intramolecular are stronger. 2.A covalent bond is 100x stronger. 3.The molecules gather together as liquids or solids at low temperatures. 4.Based on boiling points, F 2 (-188) has the weakest forces, H 2 S has the strongest (-61). 5.London forces –Are present in all compounds –Can occur between atoms or molecules –Are due to electron movement not due to  EN –Are transient in nature (dipole-dipole are more permanent ). –London forces are weaker

19. 6.A) F 2 would be lower because it is smaller. Larger atoms/molecules can have their electron clouds more easily deformed and thus have stronger London attractions and higher melting/boiling points. B) O 2 because it has only London forces. NO has a small  EN, giving it small dipoles. 7.C 8 H 18 would have the higher melting/boiling point. This is a result of the many more sites available for London forces to form. 8.1) a large  EN, 2) the small sizes of atoms. Testing concepts

20. 9.a) NH 3 : Hydrogen bonding (H + N), London. b) SF 6 : London only (it is symmetrical). c) PCl 3 :  EN=2.9-2.1. Dipole-dipole, London. d) LiCl:  EN=2.9-1.0. Ionic, (London). e) HBr:  EN=2.8-2.1. Dipole-dipole, London. f) CO 2 : London only (it is symmetrical) Challenge: In ethanol, H and O are bonded (the large  EN results in H-bonding). In dimethyl ether the O is bonded to C (a smaller  EN results in a dipole-dipole attraction rather than hydrogen bonding). Testing concepts

21. H – bonding and boiling point Q – Why are some BP suddenly high at period 2?

22. Testing concepts Boiling points increase down a group (as period increases) for two reasons: 1)  EN tends to increase 2) Size increases. A larger size means greater London forces. Boiling points are very high for H 2 O, HF, and NH 3 because these are hydrogen bonds (high  EN), creating large intermolecular forces