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Covalent bonding - electrons being shared

The octet rule can be used to rationalize why certain compounds are stable.  For instance, we used Lewis dot structures and the octet rule to rationalize the chemical formula for the compounds of oxygen, nitrogen, and carbon with hydrogen: H2O, H3N, H4C, respectively.

 

We also learned about the difference between single (one pair of shared bonding electrons), double (two pairs), and triple (three pairs) bonds. 

For instance carbon monoxide has a triple bond:

Before bonding

Single bond does not satisfy octet rule

nor does a double bond

triple bond does the trick.  Six shared "bonding" electrons and two unshared electrons for each atom satisfy the octet rule.

 

Note: when you have a molecule you do say "di", "tri" for the oxide.  For instance, carbon monoxide (CO) and carbon dioxide (CO2) vs. iron(II)chloride (FeO) and iron(III)chloride (Fe2O3).

Shapes of molecules:

We can use  Lewis structures and valence electron structures to understand the shapes of molecules. 

 

H2 - easy - linear

More complex molecules: 

Rule: electron pairs (shared bonding and unshared non-bonding) repel each other.  Molecules adjust their shape so that bonds (single, double and triple) and lone pair electrons are as far away from each other as possible. 

The number of atoms bonded to a particular atoms + the number of lone pair electrons on that atom determine its shape:

Two: linear arrangement of bonds and lone pairs

Three: trigonal planer arrangement of bonds and lone pairs

Four: tetrahedral arrangement of bonds and lone pairs

The shapes of Water H2O, Ammonia NH3 and methane CH4.  All have four pairs of valence electrons (divided among bonding and non-bonding electron pairs)

When referring to their electron pairs, they are all tetrahedral:  If we focus on the atoms only, however, we see that the molecules differ in shape.

pdb model

pdb model

pdb model

 

Bond polarity and electronegativity:

The electrons in many bonds are unevenly shared between the two atoms involved.  One atom often attracts the lion's share of the electrons.  A bond in which such uneven sharing of electrons occurs is termed a polar covalent bond.  

The polar covalent bond is an intermediate between a purely covalent bond and an ionic bond. 

The ability of an atom to attract an electron is termed its electronegativity.

Compare F and C

Carbon - Four valence electrons, each feels an effective nuclear charge of +4

Fluorine - seven valence electrons, each feels a pull of +7

fluorine's electrons are held tighter than carbons - F is more "electronegative" than carbon

electronegativity - ability of an atom to attract electrons to itself in a covalent bond.

 

Scale: no units,  F = 4.0 (Table 5.5), higher value, more electronegative

Scales like ionization energy (why?)

Bond polarity - when the electron pair is unequally distributed between two atoms - the difference in electronegativity of the atoms involved determines how equal the sharing is:

F

-

F

difference in electronegativity = 0, equal sharing, non-polar

4.0

4.0

H

-

F

difference in electronegativity = 1.9, unequal sharing, polar

2.1

4.0

If the electronegativity difference is >2.1, the bond is generally considered ionic.

HF has a polar covalent bond:

it is said to have a dipole moment

 

Polar molecules (not bonds): molecules with a net dipole moment

For more than 2 atoms, we must look at the shape to determine polarity.

H2O - OH bonds are polar covalent bonds (electronegativity difference of 1.4) as shown in black.  The sum of the two results in a net dipole moment shown in red.  Electrons are pulled towards oxygen leaving H atoms more exposed.

If H2O was linear, IT IS NOT LINEAR,  the bond dipoles would cancel and the molecule would not be polar.

DEMO- bending a stream of water.

Who cares?

Why molecules form liquids, solids, or gases,  why oil and water don't mix, depends on polarity of molecules.  Oil is non-polar, water is polar (like dissolves like)

Weak forces between molecules (holds molecular liquids and solids together)

1. Dipole interactions - attractions between polar molecules

Hydrogen bonding - special type, particularly important in water.  Gives water many of its special properties.

 

2. Dispersion forces - very weak. sloshing of electrons create temporary dipole moments.  Why non-polar molecules condense

 

Putting it all together - dissolution of a salt: