Nonpolar-covalent bond
P olar-covalent bond
δ+ δ-
Figure 9. Comparison of the electron density in (a) a non-polar, Hydrogen-Hydrogen bond and (b) polar Hydrogen-Chlorine bond
In addition to ionic and covalent bond, there is a third major type of bond - metallic bond. In solid or liquid metals, metal atoms give up electrons, as in ionic compounds. The liberated electrons however, are free to move throughout the material, rather than being held in place in negative ions.
In general, atoms of non-metals form covalent bonds with each other, atoms of metals form metallic bonds with each other, and atoms of metals form ionic bonds with atoms of non-metals. There are many exceptions, however. One common and important exception is the formation of polar-covalent bonds between metals and non-metals that do not differ greatly in electronegativity.
Sample.
Use electronegativity differences and Figure 8, to classify bonds between Sulfur and the following elements: Hydrogen, Cesium, Chlorine, Magnesium, and Oxygen. Which atom in each bond will be more negative?
The electronegativity of Sulfur is 2,5 (see Appendix 7). The more electronegative end of atom in each bond will be the atom with the larger electronegativity.
Bond from Sulfur to |
Electronegativity difference |
Bond type |
More electronegative atom |
Hydrogen |
2,5 - 2,1 = 0,4 |
Polar covalent |
Sulfur |
Cesium |
2,5 - 0,7 = 1,8 |
Ionic |
Sulfur |
Chlorine |
3,0 - 2,5 = 0,5 |
Polar covalent |
Chlorine |
Magnesium |
2,5 - 1,2 = 1,3 |
Polar covalent |
Sulfur |
Oxygen |
3,5 - 2,5 = 1,0 |
Polar covalent |
Oxygen |
Hydrogen Bonds
Hydrogen bonding differs from other uses of the word “bond” since it is a force of attraction a Hydrogen atom in one molecule and a small arom of high electronegativity (mostly Oxygen, Nitrogen, Fluorine) in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word “bond”.
When Hydrogen atoms are joined in a polar covalent bond with a small atom of high electonegativity, the partial positive charge on the Hydrogen is highly concentrated because of its small size.
Hydrogen bonding has a very important effect on the properties of water. Consider two water molecules coming close together (see Fig. 10).
Polar molecules, such as water molecules, have a weak, partial negative charge at one region of the molecule (the Oxygen atom in water) and a partial positive charge elsewhere (the Hydrogen atoms in water).
Figure 10. Formation of Hydrogen Bonds between water molecules
Thus when water molecules are close together, their positive and negative regions are attracted to the oppositely-charged regions of nearby molecules. The force of attraction, shown here as a dotted line, is a hydrogen bond. The δ+ Hydrogen is so strongly attracted to the electron pair that it is almost as if you were beginning to form a co-ordinate (donor-acceptor covalent) bond. Notice that each water molecule can potentially form four Hydrogen bonds. It doesn't go that far, but the attraction is significantly stronger than an ordinary interaction. Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water.
This is why the boiling point of water is higher than that of ammonia or Hydrogen Fluoride (see Fig. 11). In the case of Ammonia, the amount of hydrogen bonding is limited by the fact that each Nitrogen only has one electron pair. In a group of Ammonia molecules, there aren't enough electron pairs to go around to satisfy all the Hydrogens. In Hydrogen fluoride, the problem is a shortage of hydrogens. In water, there are exactly the right number of each. Water could be considered as the "perfect" Hydrogen bonded system.
The hydrogen bonds that form between water molecules account for some of the essential and unique properties of water:
The attraction created by hydrogen bonds keeps water liquid over a wider range of temperature than is found for any other molecule its size.
The energy required to break multiple hydrogen bonds causes water to have a high heat of vaporization; that is, a large amount of energy is needed to convert liquid water, where the molecules are attracted through their hydrogen bonds, to water vapor, where they are not.
Two outcomes of this:
