Q: The second law and Murphy's Law -- they're related?
A: Yes, but I shouldn't have mentioned it. First, we need to know more about bad things that can happen to us because of the behavior of chemicals in the objects around us. For a grim example, wooden houses burn down with disastrous financial loss even when people aren't killed. What's going on here in terms of energy flow?

    Wood and paper are both primarily cellulose. Paper is easier to experiment with so let's think about its burning in air. When paper catches fire and burns, there's a lot of energy given out as heat and some as yellow light. It's well known now that the products of the combustion of cellulose with the oxygen of the air are carbon dioxide and water. (The slight amount of black ash is due to the clay that was on the paper adsorbing a small amount of carbon.) Once started, the burning is spontaneous --i.e., the process goes on by itself without any further help after a match starts it -- and also the burning is really fast. Now, if all that energy is flowing out in this reaction of paper with oxygen, the paper and oxygen must have had a lot more energy inside them before the reaction than do the carbon dioxide and water after the reaction -- as shown in the diagram above.

    What's happening here is a beautiful illustration of the predictions of the second law. Systems (groups) of chemicals -- or objects made from them (like sheets of paper or houses) -- tend to react if they have more energy bound inside their molecules than do the reaction products that they can form. [[Note for Advanced students: "Energy change here means G, not just H".]] Thus, when systems react spontaneously, their reactant molecules are spreading out their internal energy in two ways: 1) part to the bonds in each molecule of the products because each of those has less energy concentrated in it than was in the molecules of the starting materials, and 2) a part to making all those product molecules move much faster (i.e., they have more kinetic energy than the original cellulose and oxygen). These faster molecules show a high temperature on a thermometer. We say they are hot, not because heat is a "something" but because heat is the process of energy transfer from one kind of matter to another -- from fast molecules of gas to the thermometer bulb or to one's hand if you get it near the flame.

Q: Wait a minute! Wait a minute!! You had to put a match to that paper to start it burning! What's spontaneous or second law about that??
A: Wait a minute yourself. Have you already forgotten that essential word "tends" in the second law?

    All the paper and wood and things made from them in the entire world tend right now to react with the oxygen in the air and form one gigantic fireball. Why don't they? Well, why don't all the mountains on earth spread out the potential energy in their high stone cliffs this second and collapse into spread out much-lower mounds of sandy particles? It's the strength of the chemical bonds (between silicon, oxygen, potassium, aluminum and other atoms and ions) that holds stone together and acts as an obstacle to the second law's immediate execution. The potential energy of high rocks/mountains is hindered from spreading out instantly.

    Just so, the strength of the chemical bonds (between carbon, hydrogen and oxygen) in cellulose holds it together and obstructs the instant spreading out of the energy inside the cellulose by reacting with the oxygen of the air. This strength prevents oxygen from instantly breaking into the cellulose molecules to form the even stronger bonds (of carbon dioxide and water) and to release large amounts of energy. However, it takes just a little extra push of energy from the hot match flame to begin to break a few quadrillion bonds in the cellulose of paper or wood so that oxygen molecules also striking those 'breaking bonds' can complete the breaks and make the first 'gazillion' CO₂ and H₂O molecules.

The initial energy push (usually from heat), shown by the small energy "hill" in the diagram below, is the activation energy, Ea, that is necessary to overcome the bond-strength obstacle to the second law in most chemical reactions. Thus, this requirement for input of an initial energy, the energy of activation, hinders both desirable and undesirable reactions from occurring. 
    An important idea is "Activation energies protect substances from change."

As these first "heated up" bonds are breaking, the oxygen from the air begins to grab carbon and hydrogen atoms to form carbon dioxide and water molecules. But the formation of new strong bonds in the CO₂ and water gives out a lot of energy -- enough to start to break many many more quadrillions of bonds of cellulose (no bond being totally broken before oxygen has simultaneously begun to form a new CO₂ and water molecule from the developing fragments). These new molecules of CO₂ and water also absorb some of the energy from the new bonds as they are formed and many move faster than twice the speed of sound. We sense those fast moving molecules as hot gas and we call it "heat".   

Q: Aha. I remember that in the Malibu fires a few years ago some houses started to blaze from the inside because heat from the nearby burning trees and brush ignited the cloth drapes inside the picture windows. Then there were others with big windows that didn't catch fire because they had aluminum blinds which were closed. That involved activation energy, right? Cotton cloth is cellulose, isn't it?
A: Yes to both questions. First of all, the glass of the windows probably got extremely hot, both from the heated air of the fire and the fire's infrared radiation. In addition, as you suggest, the intense IR radiation went right through the windows and heated the fabric drapes even more -- enough to exceed their activation energy that normally hinders their oxidation in air. They began to burn and this gave out enough energy to ignite the whole interior -- by exceeding the activation energy of oxidation of all the other flammable materials in the house. 

    Just as does every idea that we've been talking about, the concept of activation energies gives us tremendous power in understanding how the world works, even in unusual events. For instance, you've heard about the dangers of nitroglycerin, a liquid that explodes violently just from being shaken hard or jarred sharply. Do you think that its energy diagram would look like the one for cellulose above? Of course not. It must have a very low activation energy, Ea.  That leads to an extremely fast formation of hot gaseous products, an explosion (despite the  relatively smaller difference in energy between "nitro" and the products).  Explosives form hot gases so rapidly because they all have oxygen atoms as part of their molecules. Thus, those molecules don't need to wait until some molecules of atmospheric oxygen happen to hit them -- the way most substances have to do. Alfred Nobel was driven to invent a safer explosive when four workers and his brother were killed in the family nitroglycerin plant. He made what he called "dynamite" when he mixed oily nitroglycerin with some powdery silica material to form a seemingly dry solid that could be pressed into stick shape. They didn't detonate just from being hit or dropped. Obviously, therefore, an energy diagram for dynamite must look like the dotted line, a considerably higher Ea indicating that more energy must be put in, e.g., by a blasting cap, to initiate the spontaneous decomposition of the nitroglycerin. (TNT, used in armor piercing shells, is about six times more resistant to shock than nitroglycerin. Thus, you can guess at TNT's activation energy for reaction.)  Dynamite has been mainly replaced by other explosives for excavation, etc., today.

    There. We've seen some substances with low activation energies but we don't often handle nitro or TNT! How about a more important problem to many of us, rusted iron, the result of iron reacting with oxygen to form iron oxide. Of course, I'm running the risk here of opening the whole can of worms about human activity and the second law.

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