Polarity of Bonds & Molecules Bond Polarity: A polar chemical bond is one in which the electrons that are being unequally shared between two atoms. Eg. H Cl The shared electrons are actually closer to the chlorine atom since Cl is more electronegative than H (has a stronger desire for electrons).

Bond polarity can be shown by writing the symbol: d- (by the more electronegative atom in the bond) and d+ (by the less electronegative atom in the bond). Eg. Hd+--- d- Cl E.N. Cl = 3.2, H = 2.2

Predicting Bond Types The difference in electronegativity (DEN) values between two atoms in a bond indicates whether that bond is likely to be ionic or covalent. Ionic: DEN ≥ 1.7 Polar Covalent: DEN ≥ 0.5 and <1.7 Non Polar Covalent: DEN < 0.5 and > 0 Pure Covalent: DEN = 0

Eg. H—S and H—P and H—Cl all contain different types of covalent bonds E.N. H = 2.2 S = 2.6 P = 2.2 Cl = 3.2 H—S is non polar covalent (2.6 - 2.2 = 0.4) H—P is pure covalent (2.2 – 2.2 = 0) H—Cl is polar covalent (3.2 – 2.2 = 1.0)

Molecular Polarity: An entire molecule is considered to be polar if one end of the molecule is considered “negative” and the other end “positive”. The polarity of the molecule is determined by: the bonds that make up the compound as well as the number of unshared electron pairs that are around the central atom.

The “vector sum” of each bond (resultant direction of the negative and positive end) is analyzed for the entire molecule. The vector sum would cancel out in a totally symmetrical molecule, unless there are extra unshared electrons on the central atom. In that case, the most negative part of the molecule would be in the direction of those electrons. The negative part of the molecule is indicated by a tailed arrow pointing in that direction.

Relevance of Molecular Polarity It is important to know whether a molecule is polar or not since only polar molecules dissolve in polar substances like water. This is crucial in being able to carry out successful chemical reactions. Solubility Rule: “Like Dissolves Like” Polar sugar dissolves in polar water. Non polar oil dissolves in non polar gasoline. But non polar oil does not dissolve in polar water.

VSEPR Theory The atoms in a molecule like to arrange themselves such that the most electronegative atoms are as far apart as possible from each other. This minimizes the repulsion of their valence electrons. In order to do this, molecules take on basic shapes or arrangements of their atoms that depend on: Number of atoms that are around the central atom Number of extra unshared pairs of electrons around the central atom.

The theory responsible for predicting the different 3-D shapes that molecules take on is known as VSEPR (Valence Shell Electron Pair Repulsion Theory).

Molecular Shapes Attached to the central atom: 1 additional atom  linear 2 additional atoms  linear  bent (if lone pairs electrons on central atom) 3 additional atoms  planar  pyramidal(if lone pairs electrons  T-shape (if 2 lone pairs electrons on central atom)

4 additional atoms  tetrahedral  see-saw (if lone pairs electrons on central atom)  square planar (if 2 lone pairs electrons on central atom) 5 additional atoms  trigonal bipyramidal  square pyramidal (if lone pairs electrons on central atom) 6 additional atoms  octahedral

Examples of polar and non polar molecules:

Intermolecular Forces There are attractive forces between all covalent molecules, called intermolecular forces. Some are weak while others are quite strong. These forces account for the various physical properties that exist in compounds; like their state (s, l, or g), melting/boiling point, hardness/texture, solubility, and surface tension for liquids. Intramolecular forces (bonds) are very strong and hold atoms together within one molecule.

Types of Intermolecular Forces Weakest of all forces is called London dispersion which exist in all molecules; and is due to the attraction of one molecules protons to another molecules electrons, and vice versa.

Slightly stronger forces called dipole-dipole exist only in polar molecules; and is due to the attraction of one molecules negative end to another molecules positive end. Collectively London dispersion forces and dipole-dipole forces are known together as van der Waals forces.

Strongest of all intermolecular forces is hydrogen bonding which occurs only in molecules where H is directly bonded to either O, F, or N only. There is such a difference in electronegativity between these atoms that the d+H forms a relatively strong “bond” with an adjacent d-O,F,N. Many unique properties of water can be explained with hydrogen bonding; surface tension, high boiling/melting point, low density of ice