0.10 Bonding in metals and metal-non-metal salts  (Page 2/9)

In a previous study, we developed the concept of electronegativity. An atom with a high electronegativity strongly attracts the shared electrons to itself in a covalent bond. The atom with the lower electronegativity in the bond more weakly attracts the shared electrons. The result is that the bond is “polar,” meaning that one end of the bond is negatively charged and the other end is positively charged. We will assume a previous understanding of the variations of electronegativity amongst the elements. Atoms to the right in the Periodic table have higher electronegativities than those to the left. And atoms in the earlier rows of the Periodic Table have higher electronegativities than those in the later rows. Electronegativity thus generally increases from “left to right” and “down to up” in the Periodic Table. These facts will be extremely useful in understanding how and why different types and combinations of atoms form different types of bonds.

Observation 1: properties of metals

Historically, people have worked to locate, isolate, and purify metals because of their valuable properties. Most metals are both strong and malleable solids, meaning that they can be shaped, bent, pressed, flattened, and so forth without cracking or breaking. This is a very useful property. Shelters, shields, tools, and armor can be made from solids provided that they can be bent to whatever shape is desired. Since they are not brittle, metals do not break on impact so they provide excellent protection as well as excellent materials for weapons.

In the age of electricity, many metals became more valuable due to their conductivity . When a piece of metal is bridged across an electric potential, electrons flow from the negative electrode to the positive electrode, creating a current with obvious applications. By contrast, non-metals are rarely conductors and are more typically insulators. Adding to the applications of metals for electricity, metals are also ductile, meaning that they can be drawn into thin wires while maintaining strength.

And not least, most metals are actually quite attractive, with shiny, smooth, colorful finishes. This gives metals intrinsic value in addition to their usefulness. It is not surprising that gold, silver, and copper have long been used for coins and jewelry, given their beauty and their resistance to oxidation.

We can examine these properties of metals to try to understand how metal atoms are bonded together. The distinct properties of metals tells us that the bonding must be quite different from that in the covalent molecules of the non-metals we have been studying so far. These differences must relate to the differences in the properties of the individual atoms. So let’s take a look at those properties.

Perhaps the most important atomic property is, as we often have seen, the ionization energy of each metal atom. [link] shows the first ionization energy for each atom in the third and fourth rows of the Periodic Table, including both metals and non-metals. What trends do we see in these data? Two trends appear very clearly. One trend is that the ionization energies of metals are significantly lower than the ionization energies of the non-metals. Another trend is that the ionization energies of the metals do not vary much from metal to metal. This is very different from the sharp increases we see in the non-metals as we move across the periodic table.