1. 9.5-9.6 Lewis Structures for Covalent Bonding – Electronegativity and Bond Polarity

2. Single Covalent Bonds
For molecular compounds, we use Lewis structures to depict neighboring atoms as sharing some or all of their valence electrons in order to attain octets (or duets for hydrogen).
Use Lewis structures to show the bonding that occurs when a water molecule is formed. ..
H:O:H
The pair of electron between each hydrogen and central oxygen are called shared pairs, and the non-bonding pairs of electrons are called, unshared pairs. A shared pair of electrons is called a single covalent bond, and is usually drawn as a single line.
Lewis structures explain why our –gens are diatomic in nature. When two –gens bond together, they each receive complete octets (or again, a duet for the two atoms of hydrogen in a hydrogen molecule.)

3. Double and Triple Bonds
When non-metals share two pairs of electrons, a double bond is formed. Try to draw a Lewis structure of an oxygen molecule.
Double bonds are shorter and stronger than single bonds.
When two nonmetals share three pairs of electrons, a triple bond is formed. Try to draw the Lewis structure of a nitrogen molecule.
Triple bonds are even shorter and stronger than double bonds. These bonds are so strong, that they are difficult t break, making molecules with triple bonds fairly unreactive.

4. What do Lewis Structures Show us?
Lewis structures not only show us the atoms that make up the molecule, and the types of bonds that form between them, but they also explain why some elements for the compounds that they do.
They also show that covalent bonds are highly directional (these bonds exist between just two atoms sharing electrons. This sharing is usually uneven), while ionic bonds are non-directional (as you move away from the center of the ion, the forces are equal in all directions).
The uneven distribution of charge is the reason why one molecule will attract to another. This attraction called an intermolecular force, is much weaker than the attraction that hold the atoms that make up a molecule (a covalent bond which is an intramolecular force.
When a molecular compound melts, the molecules themselves stay intact, but the relatively weak intermolecular forces between the molecules are broken. This is why molecular compounds tend to have lower melting points and boiling points that ionic compounds.

5. Electronegativity and Bond Polarity
As mentioned before, electron sharing is unequal when the electrons are being shared between two different non-metals.
This is called a polar covalent bond. Example: H-F
The reason this is a polar covalent bond is because fluorine is more electronegative than hydrogen, so when the two atoms bond, F will take a greater share of the electron density.
δ+ and δ- are the symbols that are used to represent partial charges of a polar covalent bond.
Electronegativity – the ability of one atom to attract electrons to itself in a chemical bond.
A purely covalent bond (nonpolar covalent bond), is when the electrons are shared evenly between the two atoms. Ex. H-H

6. Electronegativity Trend
For the representative (main group) elements:
Electronegativity generally increases across a period.
Electronegativity decreases down a group.
Fluorine is the most electronegative element.
Francium is the least electronegative (most electropositive)
Arrange these elements in order of decreasing electronegativity:
P, Na, N, Al
N > P > Al > Na

7. Bond Polarity, Dipole Moment, and Percent Ionic Character
In order to determine how polar a bond is, an electronegativity difference (ΔEN) between the two atoms in the bond needs to be calculated.
A small ΔEN (0 – 0.39) = Purely covalent
An intermediate ΔEN ( ) = Polar covalent
A large ΔEN (2.00+) = ionic
A dipole moment (μ) occurs anytime there is a separation in positive and negative charges.
μ = q r (q = charge: unit is C (coulomb)) (r = distance: unit m (meter))
C m = D (Debye)
measured μ
Percent ionic character = x 100
μ if electrons were completely transferred = 6.2 D
A percent ionic character above 50% is said to be ionic.

8. Let’s Try a Practice Problem
Determine whether the bond formed between each pair of atoms us pure covalent, polar covalent, or ionic. a.) I and I b.) Cs and Br c.) P and O a.) pure covalent b.) ionic c.) polar covalent Let’s try another: The HCl(g) molecule has a bond length of 127 pm and a dipole moment of 1.08 D. Without doing detailed calculations, determine the best estimate for its percent ionic character. a.) 5% b.) 15% c.) 50% d.) 80% b.) 15%

9. pgs #’s 49, 52, 56 & 58
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