Covalent Bonding

-- atoms share e–

-- covalent (molecular) compounds tend to be solids with low melting points, or liquids or gases

-- three shared pairs = a triple covalent bond of e– (i.e., 6 e–)

-- two shared pairs = a double covalent bond of e– (i.e., 4 e–)

-- one shared pair = a single covalent bond of e– (i.e., 2 e–)

+ +

(–) charge density

max. EN = 4.0 (F)

bond polarity:

nonpolar covalent bond:

polar covalent bond:

electronegativity (EN): the ability of an atomin a molecule toattract e– to itself

describes the sharing of e– betweenatoms

e– shared equally

e– NOT shared equally

-- A bonded atom w/a large EN has a great ability to attract e–. -- A bonded atom w/a small EN does not attract e– very well.

min. EN = 0.7 (Cs)

-- EN values have been tabulated. (see p. 308)

nonpolar covalent

The DEN between bonded atoms approximatesthe type of bond between them.

DEN < 0.5

0.5 < DEN < 2.0

DEN > 2.0

polar covalent

ionic bond

As DEN increases,bond polarity... increases.

DEN for a C–H bond = 0.4, soall C–H bonds (such as thosein candle wax) are nonpolar.

Dipole Moments

Polar covalent molecules have a partial (–) anda partial (+) charge and are said to have adipole moment.

H–F H–F d+ d–

partialcharge O

HH

d–

d+ d+

big DEN = _____ polarity = _____ dipole momentbig big

Polar molecules tend to align themselves with eachother and with ions.

H–F H–F

H–F

H–F H–F

H–F

NO3–

NO3– NH4

+

NH4+

** Nomenclature tip: For binary compounds, the less electronegative element comes first.

-- Compounds of metals w/high ox. #’s (e.g., 4+ or higher) tend to be molecular rather than ionic.

e.g., TiO2, ZrCl4, Mn2O7

Lewis Structures (or “electron-dot structures”)

1. Sum the valence e– for all atoms. If the species is an ion, add one e– for every (–); subtract one e– for every (+).

2. Write the element symbols and connect the symbols with single bonds.

3. Complete octets for the atoms on the exterior of the structure, but NOT for H.

-- If your LS doesn’t have enough e–, place as many e–

as needed on central atom.-- If LS has too many e– OR if central atom doesn’t have an octet, use multiple bonds.

4. Count up the valence e– on your L.S. and compare that to the # from Step 1.

26 e–

Draw Lewis structures for the following species.

PCl3 P–ClCl–

Cl

.. ....

..

....

....

....

HCN 10 e– H–C–N

H–C=N

H–C–N

..

..

......

..

[ ]PO43–

P–OO–....

..

....

....

....

O

O

.. ....

32 e–3–

H2

CH3CH2OH

CO2

As the # of bonds between two atoms increases,the distance between the atoms...

2 e– H–H

20 e–

HH–C–C–O–H

H

HH

16 e–

O=C=OO–C–O ..

..

........

..

..

....

....

decreases.

This CO bond is longer (andweaker) than this CO bond.