I’ve been thinking about models and metaphors lately. Where do these show up in my field of chemistry? Well, everywhere! Why? Because in chemistry we’re trying to link tiny things we can’t see to the mesocosm that we’re used to experiencing. We try to conceptualize what we think is going on at the atomic scale by imagining objects that we can manipulate on our human scale. For example, we picture atoms as balls, and chemical bonds between atoms in a molecule as sticks connecting the balls. It’s a start, even though reality is nowhere as simple.
One crowning principle we emphasize at the G-Chem level is stability. Stable substances are those that we can macroscopically handle, touch, and taste. They stay around long enough to say something about their chemical properties, at which point we start imagining nanoscopically tiny balls and sticks of tiny things explain these properties. Now, when we say something is stable, we implicitly mean it is energetically stable. Thus, we quantify stability. As I tell my students, we don’t really know what energy is, but we can count it!
To make matters worse, something that is more stable has a lower energy, while something that is unstable is high-energy. Low is good, high is bad. That’s counter-intuitive to our other metaphors of high and low. Getting a high score on your chemistry exam is good. It’s good when your teacher thinks highly of you. In an argument or battle, you want to occupy the high ground. How you wish that the number following the dollar sign in your bank account was higher! Stable substances composed of stable molecules prefer the opposite. They want to be low. Low in energy. Sluggish. Sounds bad. But really, that’s where chemistry is headed. It’s thermodynamics!
Ah, but it gets worse! We often forget to emphasize that whenever we discuss something being stable, we mean it is relatively energetically stable to some reference state. So, if someone comes up and tells you that molecule X is stable, you should always ask: Stable relative to what? And there’s more. It will also depend on the environment your substance or molecule is in. That’s why we stress “standard conditions” in chemistry, and my students are treated to my broken record repetition of knowing what the reference state is.
What are some examples that we see in G-Chem? (Let’s assume non-changing environmental conditions of room temperature and pressure.)
Drawing “good” Lewis structures. Um, professor, what do you mean by “good”? Ah, that’s hard to say but we can use some rules-of-thumb to generate reasonable drawings of stable molecules: Follow the octet rule, don’t use large formal charges, if you do have +1 or -1 formal charges, try to match them with the less and more electronegative elements respectively, and if you can draw equivalent “resonance” structures, that helps too. (SO2 is a good example to illustrate all of this.) The notion of resonance in itself is a confusing bugbear, so let’s avoid getting into it for the moment (see Kerber, R. C. J. Chem. Educ. 2006, 83, 223-227).
After drawing some simpler structures, at some point we get to ozone. It has a beautiful symmetric triangular Lewis structure that follows all the rules. But now we throw in VSEPR Theory and argue that the internal sixty-degree bond angles are energetically bad and electron-pair repulsion (due to the Pauli Principle, I stress, not electrostatics) must be taken into account. Okay, so now we can draw bent ozone. But ozone (O3) ain’t all that stable. It decomposes into O2, the form of molecular oxygen we love to promote as life-giving. We neglect to explain why. The explanation is trickier than it looks at first glance. Many of the hand-waving ones we professors are good at trotting out are what I call “just so” explanations (after the Rudyard Kipling stories) that wither under the attack of counter-examples.
In fact, O2 isn’t all that stable, despite its Lewis structure following all the rules. Why? It depends on what else is around and its double bond is anomalously weak. Anomalous? Compared to what? Well, compared to C=C or C=O perhaps. In fact, the reason why we can burn most organic compounds for energy is because the O=O bond is weak, and a thermochemical argument can be made that it provides the lion’s share of energy released or emitted in “combusting” a fuel (see Schmidt-Rohr, K. J. Chem. Educ. 2015, 92, 2094). O2 is the “high-energy” molecule, the fuel much less so. (We’re assuming complete combustion to CO2 and H2O here.)
This leads to a distinction between kinetic and thermodynamic stability. O2 may be thermodynamically unstable compared to other common substances in the atmosphere, but you don’t see cows spontaneously combust all the time, despite their methane output. But once again, that’s relative to what else is going on and what the environmental conditions are. And I sidestepped why the O=O bond is weaker than you might expect. Clue: It’s the Pauli Principle again kicking in at shorter bond lengths. Same reason why peroxides (compounds with O–O single bonds) are reactive, one might even say unstable, although you can draw good Lewis structures with no problem.
Knowing the relative strength of chemical bonds in molecules does allow you to predict thermodynamic stability. But in a chemical reaction, bonds are broken and bonds are formed. Both processes are going on. Breaking a bond requires energy input. It is costly (energetically). But what makes the chemical reaction thermodynamically favorable? New stronger bonds are formed, and therefore the overall energy “stored” in the bonds of the chemical system is now lower. You’re going from a higher energy situation to a lower energy situation. Remember, counterintuitively, that low is good.
Students instinctively think that making lower energy compounds with stronger bonds “stores” more chemical energy, but that’s wrong – at least if you think of chemical energy as a resource you can tap when you need to convert it to electricity or heat or work. It’s the higher-energy chemical substance with weaker chemical bonds that works as your chemical “store”. This metaphor gets us into trouble because we think of a store as having more. Yes, we have more energy, but that means weaker and fewer chemical bonds where possible. If I can get this through to the students, I consider it a great achievement. It’s so challenging conceptually because the metaphors we use are counter-intuitive to our picture in everyday life of perhaps grain stored in a silo as a resource. This is part of what makes understanding chemistry challenging. We need to use metaphors. And when we use them, they’re counter-intuitive. Can’t we come up with something more intuitive? I’d like to say yes, but reality is just that much more complex.
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