Why are metal oxides generally more stable than their corresponding elemental forms?

Question

Why is the oxidation of certain metals such as iron favorable from a thermodynamic standpoint? What is it about these metal oxides that makes them inherently more thermodynamically stable than their starting, elemental forms? The only reason I can imagine is that iron, being a d-block element, has a low electronegativity and cannot stabilize valence electron density as well as oxygen, the second most electronegative element on the table.

Also, is oxidation always thermodynamically favorable?

Answer 1

The oxidation of gold is not favorable. For many other metals the metallic bonding is weaker than the bonding with oxygen.

Let's look at the steps in oxidation:

Breaking metallic bonds (endothermic).

Ionizing the metal atoms (endothermic).

Electron affinity of oxygen (exo for the first, endo for the second).

Lattice energy of the oxide (exothermic).

The lattice energy is so high that the process is favorable.

Answer 2

Both the answers above are accurate and touch on my point here. There are two factors that contribute to whether a reaction is spontaneous (if it will happen on its own) such as the one you described (iron with oxygen).

1. The thermodynamic properties of the reaction (like the net heat exchange, exothermic vs. endothermic, etc.)
2. The entropy of the reaction. This second one is often missed or not thought of. You see, the processes involved in summing the relative bond strengths, ionization energies/affinities, and lattice energies, to name a few, are examples of thermodynamic calculations that you can use to predict if a reaction will occur spontaneously. In this case, as was already stated above, the reaction between iron and oxygen is very energetically beneficent for both substances because the oxygen, in essence, ionizes the iron and forms an incredibly strong lattice. So strong it's more stable energetically than Fe itself. Another result of this stability is that the oxidized Fe is less likely than pure iron to react with atmospheric water so it's less prone to rusting, another benefit. But that's not even all! That's just number one! Remember how I mentioned entropy (that's the dispersal of energy in a system. Iron is a Dblock atom as you said. When it gives up electrons in pairs to oxygen it starts to become more stable because its Z.eff (effective nuclear charge, look it up) increases on its remaining electrons and as the oxygen gains octets, it becomes just a little more stable because its entropy (energy dispersal which is a good thing) is increasing!

The result is an increasing entropy and a decreasing internal energy for the system. These are two independent factors that often come in pairs in chemistry. But not always. The evaporation of sweat for example is endothermic and wouldn't be spontaneous is not for it's subsequent huge increase in entropy. These two principles are the circumstances that spawn spontaneous reactions and explain why iron is more stable as an oxide.

Answer 3

When you think about it, metals tend to want to lose electrons to finish with full outer shells. When they bond with oxygen, they do lose those electrons, ending in a more stable electron configuration. These ions don't go off and violently react with other things because they're ionically attracted to the oxygens they've just given their electrons to. So all their electrons are happily occupied, meaning they have no need to react and share them off. Also, because of oxygen's high electronegativity, it is unlikely to give those electrons up to another element, causing a reaction.