Re: How can two negatively valent elements combine? eg, ClO2

That's a very good question.

Using the basic model of valence, we can readily explain ionic bonds, like when chlorine "takes" one electron from a positively valent atom (like sodium) to complete its outermost electron shell and be more stable. This also lets the sodium become more stable by removing the lone electron in its outermost orbital. These two elements are now charged ions (-1 and +1 respectively), and unless dissolved in a solution, are very tightly attracted to each other.

This can also sometimes explain covalent bonding, where atoms "share" electrons to fill their outermost orbital shells. For example, an oxygen atom needs 2 electrons to complete its outermost shell. Two oxygen atoms can "share" two of their own electrons with each other. Neither oxygen "keeps" the electrons, but they are chemically bound together by the mutual sharing of electrons.

However, as you point out in your question, there are compounds that do NOT follow the rules of bonding as explained by the valence bonding theory. The main thing to understand is that ANY explanation of chemical bonding makes use of a kind of "model" or analogy that makes certain assumptions so we can more easily understand and discuss what is going on. As exceptions to the rules of a given model are found, scientists come up with newer models that can more accurately predict (or explain) all of the observed phenomena.

Molecular orbital theory builds upon the idea of individual atomic orbitals from the valence bond theory. Instead of each atom "trading" or "sharing" electrons, the atomic orbitals combine to form "molecular orbitals". In each electron shell there is a spherical "s" orbital, which holds up to 2 electrons. In slightly bigger atoms, there are "p" orbitals, which can hold 6 more electrons. Atoms at this point can be assigned two valence numbers, based on how many electrons are needed to either "fill" or "empty" the outermost shell.

In bigger atoms, there is set of orbitals available (called "d") which can hold up to 10 electrons. Normally, electrons will occupy the "lowest" orbitals possible, since it takes more energy to occupy the higher orbitals. But when enough energy is available, unusual compounds can form by making use of these higher energy orbitals. They are much less stable, and tend to break apart very easily, but they CAN form. In the example above, two oxygens bind together and are stable. But when there is an electric spark, there is enough energy available to let THREE oxygen bind together and form ozone! Ozone is unstable, but it definitely exists, even though it breaks the rules of valence bond theory.

The example you cite, chlorine dioxide, is an unstable molecule like ozone. It is VERY unstable and will naturally decompose into dioxide and dichlorine (O2 and Cl2). When chlorine dioxide decomposes, it does so explosively, even at low concentrations!

The main point here is that NONE of these theories is perfect, but each one adds to how well we understand many observed properties of molecules.

I recommend looking in a basic chemistry book at the topics of Molecular Orbital theory (MO), resonance structures and hybridization of atomic orbitals to get a better understanding of how molecules like chlorine dioxide and ozone can form.

Have a great summer!

-Jason

References:

Chemistry, 6th edition by Raymond Chang (1998).

Online fact sheet for chlorine dioxide: http://www.clo2.com/factsheet/factindex.html