0.10 Bonding in metals and metal-non-metal salts  (Page 3/9)

We have already discussed the first trend in our study of atomic structure. We explained the lower ionization energies for metals compared to non-metals from the fact that the metals have relatively lower core charges. In our previous studies, we saw that the core charge increases for atoms as we increase the atomic number in a single row of the Periodic Table. Each row consists of elements with valence electrons in the same energy shell. The metals are more “to the left” in each row, meaning that they have smaller core charges for that row. This explains the relatively lower ionization energies.

The lack of big variations in the ionization energies of metals is harder to understand. We have seen that the ionization energy of an atom is determined by the electron configuration. For the metal atoms, these electron configurations are a little trickier than for the non-metals. These are illustrated in [link] for the atoms in the fourth row of the Periodic Table, beginning with K. In each of these atoms, the 4s and 3d subshells have energies which are very close together. In K and Ca, the 4s orbital energy is lower, so the outermost electron or electrons in these two atoms are in the 4s orbital and there are no 3d electrons. For the transition metal atoms from V to Cu, these atoms have both 4s and 3d electrons and the number in each orbital depends very sensitively on exactly how many valence electrons there are and what the core charge is. As such, the actual electron configurations in [link] would have been very hard to predict and we will treat them as data which we have observed.

A clear and surprising rule in the data in [link] is that the outermost electron in each atom is always a 4s electron. As we increase the atomic number from V to Cu, there are more 3d electrons, but these are not the highest energy electrons in these atoms. Instead, these added 3d electrons increasingly shield the 4s electrons from the larger nuclear charge. The result is that there is not much increase in the core charge, so there is not much increase in the ionization energy, even with larger nuclear charge.

This model explains the data in [link] . Although this analysis probably seems complicated, it helps us to understand the important observation that all of the metals have low ionization energies.

How do these electron configurations determine the properties of metals? Or stated more specifically, how do the electrons configurations affect the bonding of metal atoms to each other, and how does this bonding determine the properties of the metals? To find out, let’s look at each property and develop a model that accounts for it. First, think about the electrical conductivity of metals. When a relatively low electric potential is applied across a piece of metal, we observe a current, which is the movement of electrons through the metal from the negative to the positive end of the electric field. The electrons in the metal respond fairly easily to that potential. For this to happen, at least some of the electrons in the metal must not be strongly attracted to their nuclei. Does this mean that they are somewhat “loose” in the metal? In fact, we have seen that this is true: the ionization energy of metal atoms is low. Perhaps the valence electrons are somewhat “loose” in the metal. A metal’s conductivity tells us that when we have many metal atoms (let’s say 1 mole, for example), there are electrons available to contribute to the current when an electric potential is applied. Thus, the valence electrons must not be localized to individual nuclei but rather are free to move about many nuclei.