In the previous page, we have learned about the two most important types of bonding: ionic bonding and covalent bonding. Both of these are ultimately driven by the desire that atoms have to be surrounded by a complete shell of electrons. They achieve this by respecetively either gaining or losing, or sharing, one or more electrons.
There are other principles which can lead to atoms bonding to each other, and we will examine here two important cases: metallic bonding, and hydrogen bonding.
Metals are well known to be solids (except for Mercury!). The bonds between metals can loosely be described as covalent bonds (due to sharing electrons), except that the metal atoms do not just share electrons with 1, 2, 3 or 4 neighbours, as in covalent bonding, but with many atoms. The structure of the metal is determined by the fact that each atom tries to be as close to as many other atoms as possible. This is shown here for one typical metal structure (adopted, for instance, by iron at some temperatures):
Can you count how many neighbours each iron atom is bonded to? Contrast this with the structure of diamond seen previously.
Because the electrons are shared with all the neighbours, it is quite easy for the electrons in metals to move around. If each "shared" electron shifts one atom to the right or left, this leads to a net shift of charge. This occurs quite easily in metals, but much less so in ionic solids, or covalent ones, where the electrons are rigidly associated with either a particular atom or ion, or a particular pair of atoms. It is because electrons can move around so easily inside metals that the latter conduct electricity.
In covalent bonds, the electrons are shared, so that each atom gets a filled shell. When the distribution of electrons in molecules is considered in detail, it becomes apparent that the "sharing" is not always perfectly "fair": often, one of the atoms gets "more" of the shared electrons than the other does.
This occurs, in particular, when atoms such as nitrogen, fluorine, or oxygen bond to hydrogen. For example, in HF (hydrogen fluoride), the structure can be described by the following "sharing" picture:
However, this structure does not tell the whole truth about the distribution of electrons in HF. Indeed, the following, "ionic" structure also respects the filled (or empty) shell rule:
In reality, HF is described by both these structures, so that the H-F bond is polar, with each atom bearing a small positive (δ+) or negative (δ-) charge. When two hydrogen fluoride molecules come close to each other, the like charges attract each other, and one gets a "molecule" of di-hydrogen fluoride as shown:
The weak "bond" between the F atom and the H is called a Hydrogen Bond, and is shown here as the dotted green line.
Hydrogen bonds can also occur between oxygen atoms and hydrogen. One of the most important types of hydrogen bonds is of this type, and is the one occurring in water. As discussed for HF, the electrons in H2O molecules are not evenly "shared": the oxygen atom has more of them than the hydrogen atoms. As a result, oxygen has a (partial) negative charge, and the hydrogens have a positive charge. When you have two water molecules close to another, a hydrogen atom on one of the molecules is attracted to the oxygen of the other molecule, to give a dimer. The structure of this dimer is shown here:
Notice how the oxygen, hydrogen, and oxygen atoms involved in the hydrogen bond are arranged more or less in a straight line. This is the preferred geometry for hydrogen bonds, and explains why only one hydrogen bond can be formed in the water dimer.
Upon going to three water molecules, it is now possible to form several hydrogen bonds. This is shown here:
How many hydrogen bonds is each water molecule involved in?
In liquid water or ice, many water molecules are close to each other, and they form dense networks of hydrogen bonds. In ice, the arrangement of the water molecules with respect to each other is regular, whereas in water, it is random. The following picture shows a typical arrangement of water molecules similar to what you might find in the liquid:
Can you see some of the hydrogen bonds? These bonds are weaker than typical covalent or ionic bonds, but nevertheless strong enough to make molecules which can hydrogen bond much more "sticky" with respect to each other than are other covalent molecules with otherwise similar properties. For example, the molecular mass of water is 18, and that of nitrogen is 28, yet nitrogen is a gas down to almost -200 degrees centigrade, whereas water is a liquid up to 100 degrees!
The cells of living things are made up of many different sorts of molecule. Two important classes of molecule are proteins and nucleic acids. In both of these molecules, parts of the (very large) molecules are involved in hydrogen bonds with other parts of the same molecules. This is very important in establishing the structure and properties of these molecules. By clicking on the link, you can view a Chime page explaining the structure of DNA, one of the most important nucleic acids, and showing the important role of hydrogen bonding.
Ionic and covalent bonding are not the only kinds of bond between atoms. Some important other types of bond include metallic bonds, and hydrogen bonds. These explain the properties of metals, e.g. that they conduct electricity, and are very important in establishing the properties of water and living cells.
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