CHEMICAL BONDS

I. Non-polar Covalent: This type of bond occurs when there is equal sharing (between the two atoms) of the electrons in the bond. Molecules such as Cl2, H2 and F2 are the usual examples.

Textbooks typically use a maximum difference of 0.2 - 0.5 electronegativity to indicate nonpolar covalent. Since textbooks vary, make sure to check with your teacher for the value he/she wants.

An example of the physical state (in room temperature) of this kind of bonds can be:

• H2 : is a colorless, odorless, tasteless, non-toxic, nonmetallic, highly combustible diatomic gas

II. Polar Covalent: This type of bond occurs when there is unequal sharing (between the two atoms) of the electrons in the bond. Molecules such as NH3 and H2O are the usual examples.

The typical rule is that bonds with an electronegativity difference less than 1.6 are considered polar. (Some textbooks or web sites use 1.7.) Obviously there is a wide range in bond polarity, with the difference in a C-Cl bond being 0.5 -- considered just barely polar -- to the difference the H-O bonds in water being 1.4 and in H-F the difference is 1.9. This last example is about as polar as a bond can get.

An example of the physical state (in room temperature) of this kind of bonds can be:

• H2O : Water is a liquid at standard ambient temperature and pressure, that transparent fluid which forms the world's streams, lakes, oceans and rain,

III. Ionic: This type of bond occurs when there is complete transfer (between the two atoms) of the electrons in the bond. Substances such as NaCl and MgCl2 are the usual examples.

The rule is that when the electronegativity difference is greater than 2.0, the bond is considered ionic.

• NaCl : Sodium chloride is the salt (solid) most responsible for the salinity of the ocean

So, let's review the rules:

1. If the electronegativity difference (usually called ΔEN) is less than 0.5, then the bond is nonpolar covalent.

2. If the ΔEN is between 0.5 and 1.6, the bond is considered polar covalent

3. If the ΔEN is greater than 2.0, then the bond is ionic.

That, of course, leaves us with a problem. What about the gap between 1.6 and 2.0? So, rule #4 is:

4. If the ΔEN is between 1.6 and 2.0 and if a metal is involved, then the bond is considered ionic. If only nonmetals are involved, the bond is considered polar covalent.

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