The modern definition of electronegativity is due to Linus Pauling. It is:
The power of an atom in a molecule to attract electrons to itself.
Pauling was able to develop a numerical scale of electronegativities. Your textbook should have a more complete list. Here are ten common elements with their values:
| F | 4.0 | C | 2.5 | |
| O | 3.5 | S | 2.5 | |
| Cl | 3.0 | H | 2.1 | |
| N | 3.0 | Na | 0.9 | |
| Br | 2.8 | K | 0.8 |
When you examine a periodic table, you will find that (excluding the noble gases) the electronegativity values tend to increase as you go to the right and up. The reverse statement is that the values tend to decrease going down and to the left. This pattern will help when you are asked to put several bonds in order from most to least ionic without using the values themselves.
Electronegativity values are useful in determining if a bond is to be classified as pure covalent, polar covalent or ionic.
What you should do is look only at the two atoms in a given bond. Calculate the difference between their electronegativity values. Only the absolute difference is important.
I. Pure 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 to indicate pure covalent. Since textbooks vary, make sure to check with your teacher for the value he/she wants. The ChemTeam will use 0.2.
One interesting example molecule is CS2. This molecule has nonpolar bonds. Sometimes a teacher will only use diatomics as examples in lecture and then spring CS2 as a test question. Since the electronegativities of C and S are both 2.5, you have a nonpolar bond.
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. Obviously there is a wide range in bond polarity, with the differences in the C-H bonds in CH4 being only 0.4 to the difference the H-O bonds in water being 1.4.
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.
So, let's review the rules:
1. If the electronegativity difference (usually called DEN) is less than 0.2, then the bond is pure covalent.
2. If the DEN is between 0.2 and 1.6, the bond is considered polar covalent
3. If the DEN is greater than 2.0, the 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 DEN 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.
So, that means compounds like HF and SiO2 are considered to be polar covalent, even though there is a large electronegativity difference.
A warning: rule #4 may not exist in your textbook. Often, the 1.6 value is used and the 1.6-2.0 range is lumped into the ionic category. So, make sure you go with what your teacher wants and do not do the "well, your're wrong because some guy on the Internet says so" thing.