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

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

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

4

C2.5

N3.0

O3.5

F4.0

P2.1

S2.5

Cl3.0Br2.8

I2.5

Some Electronegativity (E.N.) Values

H2.1

5Attention!!!!

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.

7Predicting Bond Type from E.N. (Electronegativity) ValuesConsider 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. = 2.5-2.1 =0.4 which is < 0.5.C-H is a non-polar covalent bond.

9A 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. = 3.5 -2.1 = 1.4 PolarN-H E.N. = 3.0 – 2.1 = 0.9 Polar

10An 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.ExampleNaCl E.N. = 3.0 – 0.9 = 2.1 Ionic

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

12A 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.

15Molecules may have more than Molecules may have more than one polar covalent bond.one polar covalent bond.Consider CHConsider CH22ClCl22, commonly, commonly called “methylene chloride.”called “methylene chloride.”Four electron pairs around C => Four electron pairs around C => 109° bond angle, tetrahedral 109° bond angle, tetrahedral geometry for electron pairs; geometry for electron pairs; there are also 4 atoms bonded, so there are also 4 atoms bonded, so tetrahedral molecular geometry. tetrahedral molecular geometry. (see next slide)(see next slide)

16

C

Cl

Cl HH

Faces youAway from you, behind page

andare in planeof the page

17

Picture of the 3D model of CH2Cl2

18C-H bond is non-polarC-Cl bond is polar: + - C-ClThe 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

Cl HH

-

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 CH2Cl2.Red = higher electron densityBlue = lower electron density

21This picture shows the electron density on top of the model of CH2Cl2.The orange balls are Cl and -.

22A molecule with polar bonds may be non-polar overall if the individual dipoles cancel. ClCCl4 is Cl - C - Cl Cl Each C-Cl bond is polar as +C-Cl -.

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

C

Cl

Cl ClCl

-

-

--

+

24The two net dipoles exactly cancel each other in space, so that CCl4 is a non-polar molecule, even though there are polar bonds.

Cl

C

Cl

ClCl

-

--

+

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

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

- --

+

++

molecule

27When one molecule’s electron “cloud” is distorted temporarily, it causes a neighboring molecule’s electron cloudto 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

molecule