Most atoms love to get together with other atoms. They don’t just want to get next to each other, they yearn to bond with other atoms. Not all of them do; there’s a class of elements called noble gases which are nature’s loners, their desire is to remain aloof. But there are only six of those [seven if you include Ununoctium, which doesn’t occur naturally but we’ve synthesized with nuclear reactions]; most elements really want to bond with others.
There’s a limit to how many bonds atoms like to form. Hydrogen atoms, for instance, only want to form one bond. They’re perfectly happy to bond with another hydrogen atom to form a hydrogen molecule (H2), or to bond with a chlorine atom to make hydrogen chloride (HCl). When a hydrogen atom is bonded, it’s done. No more bonds for me!
When atoms bond, they’re more stable together than alone. More stable means they release energy; to break the bond you have to give the energy back. So if you want to get energy out of chemical reactions, one way is to allow unbonded atoms to bond. What you get in return is the bond energy — to break them apart you have to put the bond energy back.
For maximum energy, you want bonds that have the most bond energy. You might have two hydrogen atoms bonded into a hydrogen molecule (H2), and two chlorine atoms bonded into a chlorine molecule (Cl2), but if you break them apart, then join each hydrogen to each chlorine, you’ll have two molecules of hydrogen chloride (HCl), a.k.a. hydrochloric acid. We can write this as a chemical reaction
H2 + Cl2 2 HCl.
You’ve still got two bonds (one in each molecule), but the bond energy of two HCl molecules is more than the bond energy of an H2 and a Cl2, so the total bond energy is greater. That means you get energy out of the reaction; however much you increase the bond energy is how much energy you get in other forms — often useful forms like light or heat.
Some atoms aren’t satisfied with just one bond; oxygen, for instance, wants two. It will bond with two hydrogen atoms to form a molecule of water (H2O). Each hydrogen has only one bond (with the oxygen), the oxygen has two bonds (one with each hydrogen), so they’re all happy in their little menage a trois. Oxygen can bond with a single other oxygen to form an oxygen molecule (O2), but to satisfy their 2-bond desire they form a double bond. Then each oxygen has two bonds, and together they’re satisfied.
Oxygen is content double-bonded to another oxygen, but it would rather single-bond to a pair of hydrogens. If you take an oxygen molecule (O2), and add two hydrogen molecules (2H2), both of the hydrogens will break up and so will the oxygen, then they’ll re-combine to form a pair of water molecules. When all is said and done, the total bond energy is more, so energy is released. In chemical reaction form it looks like this:
O2 + 2 H2 2 H2O + energy.
This process is called burning hydrogen, and the energy released can be directed to something useful (in a hydrogen-powered car, for instance) or can run amok (when hydrogen explodes). Burning is the process of bonding with oxygen.
Oxygen gets around; it really wants to bond with different types of atoms, and it’s willing to break up other bonds to get those atoms. It happens a lot; oxygen is the “homewrecker” of the molecular world, and it’s not satisfied with a single bond — it needs two.
But the real slut of the atomic world, the atom that can’t resist a molecular orgy, is carbon. Two bonds isn’t enough for carbon. It wants four.
One way to satisfy carbon is to give it four hydrogens. Each H gets a single bond (to the carbon) and the C gets four bonds (to the four hydrogens), so everybody’s happy. The result is methane (CH4), the main component of natural gas.
Then along comes a pair of oxygen molecules (2 O2). They break up into four oxygen atoms, which (homewreckers that they are) rip the CH4 apart. The four H join with two of the O to make a pair of water molecules (2 H2O). The C joins with the other two O to make carbon dioxide (CO2). The C needs four bonds, so it actually makes double bonds with each O. In reaction form,
CH4 + 2 O2 CO2 + 2 H2O + energy.
Everybody’s happy, and the bond energy is big so a lot of energy gets released. That’s why natural gas (mainly methane, CH4) makes such a high-energy fuel. In practical applications the needed oxygen is easy to come by, it’s hanging around all the time in the atmosphere.
What happens to the stuff on the right hand side of the arrow, the CO2 and 2 H2O byproducts of the reaction? We usually do the cheapest thing we can with the byproducts: dump them in the atmosphere.
The water enters the atmosphere in gaseous form, as water vapor. But water so easily changes from gas to liquid that it soon leaves the sky as rain, ice, snow, fog, or dew. The carbon dioxide, on the other hand, hangs around for a long long time. Some of it dissolves in seawater, which in a way is a “carbonated soda” (but don’t drink it). Some is taken up by plants, because they want the carbon. But only about half enters the ocean or the biosphere; the other half (a little more, actually) of the CO2 we emit stays in the air. Hundreds of years. That’s why, ever since we started burning carbon-based fuels in the industrial revolution, the amount of CO2 in the air has been going up. A lot.
It’s the Carbon
Methane isn’t the only thing we burn for fuel. We’ll burn almost anything, especially what’s cheap, convenient, powerful, and available. We started long ago with wood, then coal, then began an ever-increasing exploitation of oil and all its derivatives. These days we burn a multitude of fuels — from wood to whale oil, from propane(C3H8) in the grill to butane (C4H10) in the lighter, from octane (C8H14) in your gas tank to ethanol (C2H6O) in your gas tank. What they all have in common, including the wood, is carbon.
Carbon wants that double bond action with oxygen. And oxygen wants it just as bad. Give ’em half a chance, they’ve got a threesome going — two O for each C — and they’ll leave whatever group they’re in to do it, be it pure carbon (the essence of coal) or compounds called hydrocarbons (a bunch of C with a bunch of H) or even carbohydrates (a bunch of C with a bunch of H and some O). There’s no getting around it, carbon-based fuels will make CO2.
Most of the carbon-based fuels we use today are from ancient sources; they’re not just buried they’re fossilized. The carbon compounds in coal, oil, and natural gas are the products of plant matter being buried for millions of years. We call it fossil fuel and there’s a lot of it. We have more than enough to double, even quadruple the amount of CO2 in the atmosphere.
If you plucked a million gas molecules out of the air at random, about 400 of them would be CO2. The scientific way to say this is that the atmospheric concentration of CO2 is 400 ppmv, where “ppmv”stands for “parts per million by volume.”
It wasn’t always that way. Before the industrial revolution CO2 in the air was about 280 ppmv. We know this because in polar regions there are huge ice sheets where ice has been accumulating for many thousands of years. It begins as snow, which then compacts tighter and tighter, eventually being pressed into solid ice. As it does, bubbles of air get trapped within the ice. These days, we can drill cores through the ice, extract the trapped air bubbles, and analyze them to see what the atmosphere was like long ago. The deeper the ice, the longer ago it was laid down, the older the trapped air bubbles.
We can analyze modern air by taking a sample from any location on Earth. We’ve been doing that since 1958, when Charles Keeling started monitoring atmospheric CO2 as part of the International Geophysical Year (IGY). We’ve measured it regularly ever since, in fact today we monitor CO2 at hundreds of places around the world, but the longest record is from the Mauna Loa Atmospheric Observatory in Hawaii.
Here’s the concentration of CO2 over the last thousand years:
Old values (before 1958) are based on air trapped in the ice at Law Dome in Antarctica, new values (after 1958) are from continuous monitoring at Mauna Loa. Notice that it was steady, fluctuating only a little, for most of the last thousand years. It took a dip around the year 1600, but only a bit, some 5 or 6 ppmv.
The industrial revolution in 1750 started us on a new path. At first CO2 rose just a little, regaining the 5 or 6 ppmv reduction. About 1850 it really took off, and by 1950 was rising rapidly. Nowadays, we emit about 33 gigatonnes (billion metric tons) of CO2 into the air each year, which is enough to raise the atmospheric amount by about 2 ppmv.
Continuous monitoring since 1958 gives us a sharper image of the changes:
Regular up-and-down squiggles every year are plain to see. This happens because most of the world’s land plants are in the northern hemisphere (because most of the world’s land is in the northern hemisphere). During northern spring and summer plants grow, and they get the carbon they need by taking CO2 from the air. In autumn and winter plants decay and rot, returning it. The spring/summer withdrawal and autumn/winter return of CO2 is a visible sign of the plant world “breathing” CO2 in and out.
The important thing is that CO2 is still going up, at about 2 ppmv per year. Pre-industrial was 280, we’re now at 400, another 160 will bring us to 560, which is double the pre-industrial value. If we keep going at 2 ppmv per year, we’ll reach that milestone in another 80 years.
Carbon dioxide isn’t the only gas we’ve added to the atmosphere. Methane itself is present in small amounts; out of a million air molecules picked at random, only about two would be methane.
It didn’t used to be that way either. Before civilization itself, the concentration of methane was less than 1 ppmv. But methane is produced by the cultivation of rice, which has been going on for thousands of years. It’s also produced by cattle, which we’ve been husbanding in large numbers [contrary to popular myth, cattle don’t fart methane, they belch it]. When we mine natural gas, and when we transport it by pipeline, some of the methane leaks into the atmosphere. The result is, we’ve already more than doubled the atmosphere’s CH4.
Some gases, like chloroflourocarbons (CFCs), didn’t even exist in the atmosphere until we started emitting them. We’ve increased many atmospheric gases by a lot, and a lot of them have something in common: they’re greenhouse gases.
What, you might wonder, is a greenhouse gas? It’s a gas which makes Earth warmer by reducing how effectively it can cool itself off. How it does that is a story for another day.
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You’ve still got two bonds (one in each molecule), but the bond energy of two HCl molecules is more than the bond energy of an H2 and a Cl2, so the total bond energy is greater. That means you get energy out of the reaction;
You have this backwards here:
the bond energy of the two HCl molecules is less than the combined energy of the H2 and Cl2 molecules, the excess is released.
Phil, IIRC bond energy (aka binding energy) is measured from 0 = unbound. So every bound state is one of negative energy, and the convention of language is to drop the negative because they’re all negative. So the binding energy of the HCl is “greater”, e.g. more negative.
(Somewhat amusing that a discussion on *this* blog could turn on the convention of a baseline.)
The bond dissociation energies are H₂: 436 kJ/mol, Cl₂: 243 kJ/mol and 2HCl: 2×431 kJ/mol. So in straightforward language (presumably Tamino’s intention), the binding energy of the two HCl’s is greater than the combined energies of the H₂ and Cl₂ (by 183 kJ/mol).
But you are strictly correct, because the binding energies are really all negative — because the bound state is always a lower energy condition than the unbound state. So 2xHCl is really a lower energy condition (-2×431) than H₂ + Cl₂ (-436 + -243).
I wondered about this bit, too, so I appreciate the clarification.
Those aren’t the *binding energy* (it is positive for a bound system), those negative values are the *potential energy* of the bonds. Which is the value of the potentials of the two together and are therefore negative for a bound system.
“energy out of the reaction; however much you increase the bond energy is how much energy you get in other forms — often useful forms like light or heat.”
I realize this is a “basic” post, and you want to avoid unnecessary complications. Nevertheless, it seems to me that “useful forms like light and heat” is a possibly misleading oversimplification. First, “useful” very much depends on our point of view. From my point of view, producing changes in more complex molecules (as often occurs in biochemical reactions) is a more useful form of energy: my life depends on it.
Neglecting the energy involved in my muscular movements (heart beating, breathing, fingers typing on the keyboard, etc.) and involved in thinking (the brain is the most energy consumptive organ per unit of mass) and focusing on “light and heat” seems to direct the reader to focus on less important aspects of the use of “chemical” energy.
Very nice and lucid! Reminds me of Asimov’s essays.
[Response: High praise indeed. He’s one of my heroes.]
We’ve increased many atmospheric gases by a lot, and a lot of them have something in common: they’re greenhouse gases.
We have also significantly decreased the concentration of atmospheric oxygen.
I’d echo Barton’s “well-done”, up to and including the Asimov comparison. (I did love those science essays, among all his numerous and varied works.)
I can’t wait for Monckton or somebody to come around and call you out for slut-shaming carbon, though…
You probably have seen this data display of sea level rise… I thought it was particularly well done —
See figure #4 and others
Thanks, Richard, that’s a good piece of information. I added that as an update to an essay of mine, here:
Nice update, Doc! Perhaps ‘tireless’ is lacking a little traction where the rubber should meet the road?
I’ll add my praise for a job well done, especially since you had to cover a lot of territory in your explanation. You’re definitely a worthy successor to Isaac Asimov as you’ve become a universal teacher in the same way that he was.
Also, I think the quibbling over the sign of the energy involved in the chemical reactions was basically just nitpicking since the bottom line is that two elements combine or break apart depending on which reaction ends up in a lower energy state. Nature is basically “lazy” and doesn’t “like” to use more energy than it has to in order to do anything. So, CO2 is such a stable, long lasting molecule because such a combination of carbon and oxygen is at a very low energy state compared to any of its precursors. And of course all the extra energy contained within the original molecular bonds of O2 and any carbon compound is released to the ambient environment I thought that your explanation was just fine and it should be easy to understand by the average lay person.
I like your new two-pronged emphasis within this blog and think that your Asimovian aspect is very valuable. Keep up the great work!
OK, quibble time again. It’s not just energy. Reactions proceed to equilibrium–pressure/volume, energy/temperature and species concentration/chemical potential all play a role. Alter any of these and the equilibrium point is shifted. I only note this correction because people really should look at physical chemistry. It is one of the most beautiful achievements of physics.
I have recently done some research for my Physics class on Climate Change. I have to say that I agree with your observations. The CO2 increase in the atmosphere has caused the infrared rays that are given off by the sun to be reflected back into the Earth causing the average temperature of Earth to have increased by 1.5°F in the past century.
I have to agree with your observations. The rise in CO2 in the atmosphere is a major problem. CO2 rising in the atmosphere causes the infrared rays given off by the sun to be trapped within the atmosphere and heat up the Earth’s surface. This has caused the average temperature of the Earth’s surface to increase by 1.5°F in the past century.