Having reviewed various types of chemical bonds (part 1) and how they affect the properties of materials, we now focus on hydrogen bonds, which are important in nature, and metallic bonds, which give metals their unique properties.

Hydrogen Bonds

Hydrogen bonds are formed when a hydrogen atom is shared between two molecules. Hydrogen bonds have polarity. The hydrogen atom (+) in water is covalently attached to a very electronegative oxygen atom. It shares its partial positive charge with a second electronegative oxygen atom.

Hydrogen bonding is a strong electrostatic attraction between two independent polar molecules (i.e., molecules in which the charges are unevenly distributed, usually containing nitrogen, oxygen or fluorine). These elements have strong electron-attracting power, and the hydrogen atom serves as a bridge between them. The hydrogen bond is much weaker than the ionic or covalent bonds. When combined with chlorine, for example, hydrogen produces a chemically reactive material, an acid.

Hydrogen bonding also has a very important effect on the properties of both water and ice. Hydrogen bonding is very important in proteins and nucleic acids and, therefore, in life processes. The "unzipping" of DNA is a breaking of hydrogen bonds, which help hold the two strands of the double helix together.

Metallic Bonds

The metallic bond is responsible for the crystalline structure of pure metals. This bond cannot be ionic because all the atoms are identical. It also cannot be covalent, in the ordinary sense, because there are too few valence electrons to be shared in pairs among neighboring atoms. Instead, the valence electrons are shared collectively by all the atoms in the crystal. The electrons behave like a free gas moving within the lattice of fixed, positive ionic charges. The extreme mobility of the electrons in a metal explains its high thermal and electrical conductivity.

In the metallic bond, an atom achieves a more stable configuration by sharing the electrons in its outer shell with other atoms. Metallic bonds are typical in elements in which the valence electrons are not tightly bound with the nucleus, namely metals (thus the name metallic bonding). In this type of bond, each atom in a metal crystal contributes all the electrons in its valence shell to all other atoms in the crystal.

Another way of looking at this mechanism is to imagine that the valence electrons are not closely associated with individual atoms but instead move around among the atoms within the crystal. Therefore, the individual atoms can "slip" over one another yet remain firmly held together by the forces exerted by the electrons. This is why most metals can be hammered into thin sheets (malleable) or drawn into thin wires (ductile). The properties of metals (e.g., malleability, ductility, conductivity) tell us that their atoms possess strong bonds yet their electrons can move freely in all directions. These general observations give rise to a picture of a "positive ions in a sea of electrons" used to describe metallic bonding.

The relationship between the thermal conductivity of elements and their location on the periodical table is shown in Figure 1. Copper, silver and gold – the most thermally conductive metals – are located in the center.