1 New subject... Why does stuff stick together? In covalent bonds electrons are not always shared “equally.” Some nuclei have a stronger pull for electrons than others. Electronegativity Electronegativity describes the relative strength of attraction an atom has for e -.

2 Metalslow electronegativity Metals have very low electronegativity values -- they give up e - easily. Non-metalslarger electronegativity Non-metals have larger electronegativity values -- they attract e -. across a periodgoing up a Group or column. Electronegativity increases: across a period, & going up a Group or column.

3 Most electronegative element = Fluorine (F). Least Electronegative = Cesium (Cs)/Francium (Fr). Pauling electronegativity scale uses rating. Most electronegative = 4 Not electronegative = 0

4 C 2.5 N 3.0 O 3.5 F 4.0 P 2.1 S 2.5 Cl 3.0 Br 2.8 I 2.5 Some Electronegativity (E.N.) Values H 2.1

5 Attention!!!! Do NOT memorize the electronegativity values – they are provided on tests. Use the table of E.N. values in the textbook for homework. Attention!!!!

6 When two non-metals form a covalent bond, the e - pair may not be evenly shared between the two nuclei. If one element is more electronegative than the other, the e - pair is pulled closer to that nucleus.

7 Predicting Bond Type from E.N. (Electronegativity) Values Consider a bond between atoms A and B, A-B. Calculate the difference in the E.N. values (or  E.N.) for the two elements A & B = Larger electronegativity – smaller electronegativity

8 A non-polar covalent bond has  E.N. less than 0.5, or in English: when the difference in electro- negativity values is less than 0.5, we call the bond non-polar covalent. Example C-H bond  E.N. = =0.4 which is < 0.5. C-H is a non-polar covalent bond.

9 A polar covalent bond has  E.N. between 0.5 and 1.9. Or, when the difference in electronegativity values is 0.5 or more up to 1.9, then we call the bond polar covalent. Examples O-H  E.N. = = 1.4 Polar N-H  E.N. = 3.0 – 2.1 = 0.9 Polar

10 An ionic bond has  E.N. greater than 1.9, or when the difference in electronegativity values is more than 1.9, then we call the bond ionic. Recall that a metal + non-metal also signifies an ionic bond. Example NaCl  E.N. = 3.0 – 0.9 = 2.1 Ionic

11 Another Example Cl is more EN than H: H - ClCl is more EN than H: the electron pair in the bond is pulled closer to Cl. The H-Cl bond is a polar covalent bond, shown as: The H-Cl bond is a polar covalent bond, shown as:  + H-Cl   The symbol  + or  - means “partial - charge.”

12 A polar covalent bond has a “dipole.” (  + and  - ends). Molecules like HCl that have only one bond associate or stick together by dipole-dipole interactions: the attractions between areas of opposite charge density in molecules. See illustration on next slide….

13  +  -  +  - H-Cl …… H-Cl …… H-Cl …… H-Cl Cl-H …… Cl-H …… Cl-H …… Cl-H  -  +  -  + : : : : : : : : Dotted lines = intermolecular attraction

14 The intermolecular forces are far weaker than the strength of the covalent bond between H and Cl (or any covalently bonded atoms). But, the extra small attractions are enough to explain why the molecules “associate,” “cluster,” or “stick” together.

15 Molecules may have more than one polar covalent bond. Consider CH 2 Cl 2, commonlycalled “methylene chloride.” Consider CH 2 Cl 2, commonly called “methylene chloride.” Four electron pairs around C => 109° bond angle, tetrahedral geometry for electron pairs; there are also 4 atoms bonded, so tetrahedral molecular geometry. (see next slide)

16 C Cl H H Faces you Away from you, behind page and are in plane of the page

17 Picture of the 3D model of CH 2 Cl 2

18 C-H bond is non-polar C-Cl bond is polar:  +  - C-Cl The two dipoles in the 2 C-Cl bonds point in different directions. The net or overall dipole acts like it is pointing as shown on next slide:

19 -- ++ C Cl H H -- Overall electron density pulled in this direction Imagine the molecule takes up a roughly spherical volume of space, one side of the sphere is electron rich, the other is electron poor:

20 A plot of electron density in CH 2 Cl 2. Red = higher electron density Blue = lower electron density

21 This picture shows the electron density on top of the model of CH 2 Cl 2.The orange balls are Cl and  -.

22 A molecule with polar bonds may be non-polar overall if the individual dipoles cancel. Cl CCl 4 is Cl - C - Cl Cl Each C-Cl bond is polar as  + C-Cl  -.

23 But each C-Cl bond points is 109° to the other 3 C-Cl bonds. C Cl -- -- -- -- ++

24 The two net dipoles exactly cancel each other in space, so that CCl 4 is a non-polar molecule, even though there are polar bonds. Cl C -- -- -- ++

25 Non-polar molecules attract each other by forming brief, weak inter- molecular forces called London dispersion forces. These result from instantaneous and Short-lived imbalances of valence electrons. For those brief moments (on molecular scale), the molecule has regions of higher and lower e - density.

26 Roughly spherical volume of non-polar molecule; instantaneous dipole created by uneven distribution of e - density: -- -- -- ++ ++ ++ molecule

27 When one molecule’s electron “cloud” is distorted temporarily, it causes a neighboring molecule’s electron cloud to reorient itself in an opposite way. This causes a very weak, very brief (but very real!) electrostatic attraction between the two molecules:

28 -- -- ++ ++ ++ -- -- ++ ++ ++ -- -- = London dispersion force; weak, instantaneous attraction molecule