Chapter 2: Polar Covalent Bonds; Acids and Bases

If the electron pairs in covalent bonds were donated and shared absolutely evenly there would be no fixed local charges within a molecule. Although this is true for diatomic elements such as H2, N2 and O2, most covalent compounds show some degree of local charge separation, resulting in bond and / or molecular dipoles.

Polar Covalent Bonds

H
2.2

Electronegativity Values
for Some Elements

Li
1.0

Be
1.5

B
2.0

C
2.5

N
3.1

O
3.5

F
4.1

Na
0.9

Mg
1.2

Al
1.5

Si
1.7

P
2.1

S
2.4

Cl
2.8

K
0.8

Ca
1.0

Ga
1.8

Ge
2.0

As
2.2

Se
2.5

Br
2.7

Polarity is a separation of charge due to unequal sharing of electrons. The more electronegative element has a partial negative charge. The less electronegative element has a partial positive charge.

 

Because of their differing nuclear charges, and as a result of shielding by inner electron shells, the different atoms of the periodic table have different affinities for nearby electrons. The ability of an element to attract or hold onto electrons is called electronegativity. A rough quantitative scale of electronegativity values has been established, and some of these are given in the table to the right. A larger number on this scale signifies a greater affinity for electrons. Fluorine has the greatest electronegativity of all the elements, and the heavier alkali metals such as potassium rubidium and cesium have the lowest electronegativities. It should be noted that carbon is about in the middle of the electronegativity range, and is slightly more electronegative than hydrogen.
When two different atoms are bonded covalently, the shared electrons are attracted to the more electronegative atom of the bond, resulting in a shift of electron density toward the more electronegative atom. Such a covalent bond is polar, and will have a dipole moment (one end is positive and the other end negative). The degree of polarity and the magnitude of the bond dipole will be proportional to the difference in electronegativity of the bonded atoms. Thus a O-H bond is more polar than a C-H bond, with the hydrogen atom of the former being more positive than the hydrogen bonded to carbon. Likewise, C-Cl and C-Li bonds are both polar, but the carbon end is positive in the former and negative in the latter. The dipolar nature of these bonds is often indicated by a partial charge notation (d+/-) or by an arrow pointing to the negative end of the bond.

Although there is a small electronegativity difference between carbon and hydrogen, the C-H bond is regarded as weakly polar at best, and hydrocarbons in general are considered to be non-polar compounds.

Molecular polarity depends on presence of polar bonds, whether the dipole moments cancel, and the symmetry of the molecule.

 

Formal Charges


Smallest formal charge for atoms is preferred.

Most electronegative element gets negative formal charge.

Least electronegative element gets positive formal charge.

 

 

A large local charge separation usually results when a shared electron pair is donated unilaterally. The three Kekulé formulas shown here illustrate this condition.

In the formula for ozone the central oxygen atom has three bonds and a full positive charge while the right hand oxygen has a single bond and is negatively charged. The overall charge of the ozone molecule is therefore zero. Similarly, nitromethane has a positive-charged nitrogen and a negative-charged oxygen, the total molecular charge again being zero. Finally, azide anion has two negative-charged nitrogens and one positive-charged nitrogen, the total charge being minus one.
In general, for covalently bonded atoms having valence shell electron octets, if the number of covalent bonds to an atom is greater than its normal valence it will carry a positive charge. If the number of covalent bonds to an atom is less than its normal valence it will carry a negative charge. The formal charge on an atom may also be calculated by the following formula:

Formal
Charge

=

Valence Electrons
in Neutral Atom

-

[

Unshared Valence
Electrons

+

Half the Shared
Electrons

]

 

Resonance

 

Resonance will occur when 3 or more p-orbitals overlap to form a molecular orbital

Resonance allows electrons to spread out over a larger region of space

Resonance therefore leads to a lowering of energy

Reactions leading to resonance are favored

 

 

Rules for Resonance Forms

 

1                    Individual resonance forms are imaginary, not real.

a.       Real structure is a hybrid of the different resonance forms.

b.      Molecular structure is not changed between resonance forms.

c.       Only difference is how we represent the resonance forms on paper.

2                    Resonance forms differ only in the placement of p or nonbonding electrons.

a.       Position and hybridization of atoms do not change from one resonance structure to another.

b.      Curved arrows indicate the movement of electrons.

3                    Different resonance forms of a substance don’t have to be equivalent.

a.       More stable resonance structures contribute more to the hybrid than less stable resonance structures.

b.      Negative charges will be preferentially on more electronegative elements.

c.       Positive charges will on less electronegative elements.

4                    Resonance forms must be valid Lewis structures and obey normal rules of valency.

a.       Octet rule still applies.

5                    The resonance hybrid is more stable than any individual resonance form.

a.       The possibility of resonance leads to more stability since electrons are not localized.

 

Acids and Bases: The Brønsted—Lowry Definition

Acids donate a proton (H+)

Bases accept a proton

The strength of acids and bases is relative to their ability to loose or gain a proton

Loosing a proton creates a conjugate base

Gaining a proton creates a conjugate acid

 

 

Acid and Base Strength

 

The strength of an acid or base can be determined by the pKa value.

The magnitude of  the pKa  is related to the relative electron density around the hydrogen, and the relative stability of the resulting ion.

 

 

 

 

 

 

 

Predicting Acid—Base Reactions from pKa Values

It is possible to predict the direction of an acid base reaction based on the relative strength of the acid and base and their conjugates. The reaction to the right demonstrates the trend that strong proceeds toward weak.

 

 

 

 

Organic Acids and Organic Bases

 

From the Brønsted—Lowry perspective, an organic acid looses a proton. Very weak organic acids (large pKa values) require strong bases to abstract a hydrogen, and stronger organic acids (small pKa values) require weaker bases to abstract the hydrogen. Generally, the stronger organic acids contain a carboxylic acid group –COOH. The ion formed from the loss of the proton leads to a stable resonance structure with the negative charge distributed in the p-network of the OCO group.

 

Organic bases are formed when an organic acid looses hydrogen. These conjugate bases are often quite strong. The stronger the organic base, the larger the electron density at the basic site. Organic bases are best looked at from the Lewis perspective.

 

Acids and Bases: The Lewis Definition

 

Lewis Bases donate electrons.

Lewis Acids accept electrons.

The strength of acids and bases is relative to their ability to gain or loose electrons.

Electron rich areas are basic.

Electron poor areas are acidic.

 

 

BF3 + CH3OCH3                BF3O(CH3)2

 

 

 

 

 

 

 

Drawing Chemical Structures

 

CH3CH3

 

CH3CH­2CH3

 

CH3CH­2CH­2CH3