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Polar Covalent Bonds



H
2.20

Electronegativity Values
for Some Elements


Li
0.98

Be
1.57

B
2.04

C
2.55

N
3.04

O
3.44

F
3.98

Na
0.90

Mg
1.31

Al
1.61

Si
1.90

P
2.19

S
2.58

Cl
3.16

K
0.82

Ca
1.00

Ga
1.81

Ge
2.01

As
2.18

Se
2.55

Br
2.96



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 was established by Linus Pauling, 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 (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 (δ+/–) 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.

The shift of electron density in a covalent bond toward the more electronegative atom or group can be observed in several ways. For bonds to hydrogen, acidity is one criterion. If the bonding electron pair moves away from the hydrogen nucleus the proton will be more easily transfered to a base (it will be more acidic). A comparison of the acidities of methane, water and hydrofluoric acid is instructive. Methane is essentially non-acidic, since the C–H bond is nearly non-polar. As noted above, the O–H bond of water is polar, and it is at least 25 powers of ten more acidic than methane. H–F is over 12 powers of ten more acidic than water as a consequence of the greater electronegativity difference in its atoms.
Electronegativity differences may be transmitted through connecting covalent bonds by an inductive effect. Replacing one of the hydrogens of water by a more electronegative atom increases the acidity of the remaining O–H bond. Thus hydrogen peroxide, HO–O–H, is ten thousand times more acidic than water, and hypochlorous acid, Cl–O–H is one hundred million times more acidic. This inductive transfer of polarity tapers off as the number of transmitting bonds increases, and the presence of more than one highly electronegative atom has a cumulative effect. For example, trifluoro ethanol, CF3CH2–O–H is about ten thousand times more acidic than ethanol, CH3CH2–O–H.
Classification of Chemical Bond Types
chemical bond represents the net attraction that keeps atoms near each other in most material samples.  It is a consequence of the electrical attraction between oppositely charged particles in atoms--namely electrons and protons.

Because there exists a large number and a diverse arrangement of electrons and protons in the various atoms of most substances, a precise understanding of all the complex electrical interactions can be challenging.  However, some simplified models of these interactions allow us to predict many important properties.  



First, we divide bonds up into two major categories: primary bonds and secondary bonds.  Primary bonds are the strong bonds between the tightly clustered atoms that give any pure substance its characteristic properties.  Secondary bonds (also known as interparticleintermolecular, or Van der Waals attractions) are the relatively weaker attractions between nearby atoms or molecules that are important in most liquids (especially liquid mixtures) and some solids.


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