Topic 7.9: Introduction to Le Châtelier’s Principle

Le Chatelier’s principle: the ammonia machine at Oppau

Every grain on this wheat plant is built partly out of nitrogen, and so is every protein in your body, and the strange thing is that you’re breathing a sea of it right now. Almost four out of every five molecules in the air are nitrogen, two nitrogen atoms locked together by a triple bond so strong that a wheat plant can’t use them at all.

For most of history, farmers got extra nitrogen from manure, and later from saltpeter dug out of the desert in Chile. Then in September 1913, at Oppau on the Rhine in Germany, the chemical company BASF switched on the first factory that made ammonia straight out of the nitrogen in the air.

Let’s go inside. In here, in the dark, steel converters squeezed nitrogen and hydrogen together at about two hundred times the pressure of the air around you, and heated them to around five hundred degrees Celsius. Out comes ammonia, N H three, one nitrogen holding on to three hydrogens, and ammonia is what nitrogen fertilizer is made from.

Researchers estimated that by 2008, fertilizer made this way was feeding almost half the people on Earth, which means there’s a good chance some of the nitrogen in your own food went through a machine like this one. But it took a fight to get here, because this reaction runs both ways at once, so ammonia keeps falling apart even while it forms, and you always end up with a mixture.

Two famous chemists tried to make this reaction work and gave up. And one of them is the man whose name is on today’s lesson. So here’s the machine. It has a gauge that shows how much ammonia is in the gas, and a lever for how hot the converter runs. Heat speeds reactions up, so if I push the converter hotter and let it settle, what do you think happens to the amount of ammonia in the gas?

Here we go, hotter, and the ammonia goes down. The heat made it worse, but without heat this reaction barely goes at all, and that trap is what stumped chemists for years. By the end of this lesson, you’ll be able to run this machine. You’ll know which handles squeeze out more ammonia. And you’ll know which ones take it away.

And you’ll be able to prove, with numbers, which way the reaction will go before it moves at all. That’s exactly what the AP exam will ask you to do. And to do any of that, we have to go inside the converter. All the way down, into the gas itself.

Push on it, and it pushes back

Down between the grains of the catalyst, here’s a small sample of that gas, drawn as a model, with nitrogen in blue pairs, hydrogen in little white pairs, and ammonia, one blue nitrogen holding three white hydrogens. Each little flash is a reaction, and they go both ways, some making ammonia and some breaking it apart.

This is equilibrium, with the forward and reverse reactions running at the same rate, so the amounts hold steady and these three lines stay flat, even though nothing in there ever stops. So let’s push on it. I’ll open a valve and pour in extra nitrogen, and you can see the blue line jump straight up.

Now watch what the gas does about it. Nitrogen and hydrogen start drifting down while ammonia climbs, because the extra nitrogen speeds up the forward reaction for a while. Let’s freeze three moments side by side, before I touched it, right after I poured the nitrogen in, and once it settled down again. Look at the nitrogen in the last picture. The gas used up some of what I added, but not all of it, and it ends up holding more ammonia than before.

This is the pattern behind everything we’ll do today. When you disturb a system at equilibrium, it shifts in the direction that partly undoes the disturbance. And notice the word partly. The system never gets all the way back to where it started. In the model’s numbers, nitrogen jumped from zero point three zero atmospheres to zero point six zero when I poured it in, and then it settled at about zero point five eight, nowhere near back down at zero point three zero.

And ammonia went from zero point two zero up to about zero point two four, which is the shift you just watched. This is called Le Chatelier’s principle, after the French chemist Henry Le Chatelier, and he’s the man I promised you at the start. In 1901 he tried to make ammonia from nitrogen and hydrogen himself, at high pressure and high temperature, but air got into his gas, his apparatus exploded, and he gave the whole project up.

Then a German chemist named Fritz Haber found only a trace of ammonia at ordinary pressure, and he later wrote that by 1905 it seemed useless to pursue the problem any further. He left it alone for about three years, and then he came back to it, this time at much higher pressure. So the man whose principle explains this whole factory gave up on this very reaction, and when Haber came back, what finally made it work was pushing the equilibrium exactly the way that principle predicts.

Let’s try the opposite kind of push. This time I’ll open an outlet and pull half of the ammonia out of the gas. The ammonia line drops, then climbs partway back, because the gas makes more ammonia to replace some of what I took.

And that’s how ammonia plants work. The gas leaving the converter is cooled until the ammonia turns into a clear liquid and drips out, and the leftover nitrogen and hydrogen go around again for another try. The gas leaving Haber’s own converter was only about five percent ammonia, and he wrote that once the gas went around and around, that five percent was enough to turn his experiment into a means of manufacture.

Here’s a graph of that same gas, the kind the AP exam loves, and at the time marked t one, the operator did something to it. Look at which line jumps straight up at t one, and how the other two respond afterward. What did the operator do?

Now you try it in the gas. Without changing the temperature or the volume, get this gas to hold more ammonia than it has right now.

Squeeze it

Now for the biggest handle in this factory, the pressure. Count the molecules of gas on each side of the equation. On the left there’s one nitrogen and three hydrogens, so four molecules of gas, and on the right there are just two molecules of ammonia. So every time the reaction runs forward, four molecules of gas become two, and fewer molecules push less on the walls.

Now I’ll squeeze the gas into half the space, and all three lines jump up, because the same molecules are crammed into half the room. Then ammonia keeps climbing while nitrogen and hydrogen fall, because turning four molecules into two takes away some of that extra pressure. So squeezing an equilibrium of gases shifts it toward the side with fewer gas molecules, and for this reaction that’s the ammonia side.

So when Haber came back to the problem in 1908, he went for the highest pressure he could get, with a compressor that could reach two hundred atmospheres, because squeezing that hard is what made the ammonia worth chasing. Squeezing this exact reaction is one of the sample questions the AP course publishes, so a little later we’re going to come back and solve it with numbers.

You can watch a squeeze with your own eyes in a different gas. This sealed syringe holds brown nitrogen dioxide and a colorless gas called dinitrogen tetroxide, at equilibrium. One colorless molecule splits into two brown ones, and two brown ones can join back into one, so the color tells you how much brown gas there is. Watch what happens when I push the plunger in. It gets darker instantly, because the same brown molecules are squeezed into half the space, and then, as I hold it there, it fades.

Here are the three moments side by side. Before, right after the squeeze, much darker, with the brown gas twice as concentrated, and after it settles, paler than that, but still darker than where it started. That fade is the gas pushing back, with pairs of brown molecules joining into colorless ones, because fewer molecules push less on the walls.

Now I’m going to pull the plunger out to double the volume, and right after I pull, the gas will look paler. As it settles after that, what happens to the color? Here goes. Paler right away, and then it browns up a little as it settles, though it never gets back to the color it started with.

Heat is part of the equation

Now let’s figure out why the heat made the ammonia go down. Making ammonia gives off heat, about ninety-two kilojoules for every two moles of ammonia, so the reaction is exothermic. So I can write the heat right into the equation. It goes on the product side. For this reaction, heat acts just like a product.

So heating the converter is like adding a product, and the system pushes back by running in reverse, using up some of that heat and turning ammonia back into nitrogen and hydrogen. Cooling it does the opposite and pushes the reaction toward more ammonia. The brown gas works the other way around. Splitting one colorless molecule into two brown ones takes in heat, so for that reaction heat goes on the left, like a reactant.

So if I dip the syringe into hot water, the brown gets deeper, because adding heat pushes the reaction toward the brown side. And in ice water most of the brown fades away, because taking heat away pushes the reaction toward the colorless side. Now you try the baths. The plunger is clamped, so you can’t change the volume. Using only the baths, make the gas in the syringe as pale as you can.

Up to now, heating the gas made it shift. A shift can work the other way, too, and change the temperature. Pour extra nitrogen into the converter, and the gas shifts toward ammonia. Making ammonia gives off heat. In an insulated converter that heat stays in the gas, and the gas gets hotter.

Back to the heat lever. At equilibrium, colder means more ammonia. But the gas only spends a short time in the converter. When it’s cold, the reaction barely gets going, so little ammonia comes out.

Haber needed a catalyst that worked fast at five or six hundred degrees. The plant’s iron catalyst gets the gas to equilibrium faster, but look where the lines end up, in exactly the same place as without it.

A catalyst makes the reaction faster, not further, so the plant runs hot to get speed and squeezes hard to win back the ammonia that the heat costs it. For the ammonia reaction at equilibrium, which one of these changes would raise the amount of ammonia once it settles?

Color, pH, and a push that does nothing

pH can tell us which way an equilibrium moved, too. This flask holds some of the factory’s ammonia, dissolved in water. A little of the ammonia grabs a hydrogen ion from water, making ammonium ions and hydroxide ions, and those hydroxide ions make the solution basic, with a high pH. The indicator in here, phenolphthalein, turns pink when there are plenty of hydroxide ions, and I’m about to tip in ammonium chloride, which adds a lot more ammonium ions.

Before I tip it in, make a prediction. When the ammonium chloride dissolves, what will happen to the pink color and to the pH? Here it goes. Look at the pink disappearing around the crystals as they sink. Extra ammonium pushes the reaction backward and uses up hydroxide ions, so the pH drops and the pink fades.

The simplest push on a solution is adding water. Diluting works like giving a gas more room. The equilibrium shifts toward the side with more dissolved particles. Our ammonia solution has one dissolved particle on the left. Water is the solvent, so it doesn’t count.

The right side has two. Diluting it shifts the reaction forward. Watch the pH as I double the volume with water. The water halves the hydroxide right away. On its own, that drops the pH to about ten point eight.

Then the forward shift makes some of that hydroxide back. The pH climbs to about eleven and stays there. Partly undone, like the syringe. Quick check on the hydroxide concentration. After the shift settles, how does it compare with where it started?

Back in the factory, I’ll lock the converter so its volume can’t change, and pump in helium, which doesn’t react with anything in here. Watch the pressure gauge climb. So with more pressure and the volume locked, which way do you think the reaction shifts?

Look at the three lines, because they don’t move at all. The pressure went up and the reaction didn’t care, so we need a sharper tool that tells us exactly when a push counts.