CHEM 125a: Freshman Organic Chemistry I

Lecture 20

 - Rise of the Atomic Theory (1790-1805)


This lecture traces the development of elemental analysis as a technique for the determination of the composition of organic compounds beginning with Lavoisier’s early combustion and fermentation experiments, which showed a new, if naïve, attitude toward handling experimental data. Dalton’s atomic theory was consistent with the empirical laws of definite, equivalent, and multiple proportions. The basis of our current notation and of precise analysis was established by Berzelius, but confusion about atomic weight multiples, which could have been clarified early by the law of Avogadro and Gay-Lussac, would persist for more than half a century.

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Freshman Organic Chemistry I

CHEM 125a - Lecture 20 - Rise of the Atomic Theory (1790-1805)

Chapter 1. The Development of Elemental Analysis: Lavoisier’s Early Combustion and Fermentation Experiments [00:00:00]

Professor Michael McBride: Okay, as you may remember, way back before the exam, we’d started looking at how things really happened, how people were able to figure out about bonds, and how atoms were arranged, and molecules reacted, before there were the powerful techniques that we — that developed mostly in the last twenty-five or thirty years, as far as their practical application in organic chemistry; but how they found these out before that time. And it started with what we call the “Chemical Revolution”, which was launched by Lavoisier, with this book, in 1789, the same year as the French Revolution, The Elementary Treatise of Chemistry. Which, you remember, he started work on just because he was interested in improving nomenclature, but found that there was such a tight coupling between nomenclature and the science — that is, the facts and the theory — that he couldn’t improve any one of them without improving the others. So we saw last time that he developed a — that in the process of developing a new nomenclature, he also developed instruments, so that he could weigh gases, because weighing turns out to be such an important feature in all of nineteenth century chemistry, no doubt the most important single technique. But he also could measure heat, which didn’t have any weight. For that he developed a calorimeter together with Laplace, and we saw how that worked by melting ice. So now we’re going to look at some of the other developments from this book, in particular the idea of oxidation.

There’s no doubt that the most important element for Lavoisier was oxygen, because of its key role in his chemistry, so that all other elements were called bases, or radicals. Right? And they could react with oxygen to become new things. There were several degrees of oxidation, as shown in this table from the English translation of his book. So first you get an oxide, oxides of all the elements that are listed in the first table. For example, hydrogen when oxidized would give water; that’s the second row. But the first row is caloric. This was sort of a stretch, at least it seems so from our perspective, that caloric, heat, when it gets oxidized, yields oxygen gas. But then there were further degrees of oxidation if you reacted with more oxygen. So the second degree of oxidation, according to his theory, was to give an acid, that the suffix was to be systematically o-u-s, so -ous. And those names we use now, like cuprous, ferrous, come directly from this work of Lavoisier. So that’s the lower oxidation state of an acid. Right? So you have, for example, nitrous acid, carbonous acid, sulfurous acid, phosphorous acid, he said. And then you have a still higher degree of oxidation, the third degree, which is the -ic acid. So you have nitric acid, carbonic acid, sulfuric acid. And notice that many of these were not known; and many of them still are not known because the theory wasn’t exactly right. For example, carbonous acid was not known at that time; you see in the second entry in the second row, second column. Okay? And then there was still a fourth degree of oxidation, which he called the oxygenated ic acid. He didn’t use the peroxy acid, as we would often use nowadays. But he used, for example, oxygenated nitric acid, which was unknown. Oxygenated carbonic acid was unknown. Same for sulfuric or phosphoric. And oxygenated muriatic acid was — used to be called dephlogisticated marine acid. But he has a more systematic name, showing its degree of oxidation. So these -ous and -ic names, that we still use, come directly from Lavoisier.

And probably the next most important tool he developed was one for doing elemental analysis. Remember, you could do proximate analysis, what other compounds are present in the substance you’re studying. Or you could do ultimate analysis and find out how much of the elements there was in each. And this was how to do elemental analysis of an organic oil. So it has carbon, hydrogen, oxygen in it. Okay? So this big upside-down barrel, inside another barrel with water in it, is how he got air for the burning, because the upside-down barrel would be lowered into the water, which would create pressure in the oxygen that’s trapped inside. It would then flow out through the pipe you see coming from the tank, through a dryer and into this big, vertical cylinder where the combustion takes place, fed by an oil supply which is siphoned in as needed. So there’s a lamp inside, burning the oil with the air, and the vapors generated by the combustion come up and out the tube to the right. So there’s the lamp. The gases come out and water gets condensed in the spiral tube, in here, and drips down and gets collected in the jar that you see. And the part that doesn’t condense and get collected into the jar, gets absorbed in that horizontal tube that has a calcium chloride drying agent in it. So at the end of the experiment you weigh how much water there is in the jar and how much the calcium chloride has increased in weight, and that tells you how much water there was.

And what does that tell you about the compound? How much what? What element is being measured in the oil, by weighing the water? It’s measuring the hydrogen. So he had to know what percent, by weight, water is, of hydrogen. Then if he measured the water that came from combustion, he measured how much hydrogen there was in the oil sample that was being burned. And then the gases continued to go through limewater, calcium hydroxide. So that would absorb CO2, as calcium carbonate. So those would increase in mass. And he had two of them, so whatever got through the first one would hopefully be collected in the next one. And then there’s another drying tube, on further from that, because these solutions, as you pass gas over them, they lose a little bit of their water. Right? So they’d be losing weight, as well as gaining weight from the CO2 that’s being absorbed. Right? But that water that’s lost is reabsorbed in that final tube. So you weigh it at the end, and now you have — you’ve corrected for whatever loss of water there was from the bulbs, and now you know what weight of CO2. And if know what fraction of CO2 is carbon — and there’s a homework problem for Friday that you can look at. You can work in groups on that, that shows you how he knew how much — what weight percentage of CO2 was carbon. So having then measured how much CO2, he knew how much carbon there was in the oil. So he measured the hydrogen, he measured the carbon. So that’s fine if it’s a hydrocarbon oil. How about if there was oxygen in the oil? Notice, he’s adding all kinds of oxygen from the air. So how’s he going to know how much oxygen there was, if it was an alcohol or something that he was burning?

Student: [inaudible]

Professor Michael McBride: Sam? Can’t hear you.

Student: The leftovers.

Professor Michael McBride: It’s the leftovers. You see how much oil got burned, how much carbon, how much hydrogen, and you assume there’s nothing else there, except oxygen. So you get oxygen by difference. And that’s the way oxygen is almost always measured, is by difference. Right? So you assume — and you have to be sure there are no other elements there, if you want that to be true, obviously. Okay, so this was a big operation and it took, you know, somebody to control the air supply to make sure it was constant. Notice, incidentally, there’s a T here. So there was another one of these things over here, another one of these double barrels. Right? Why?

Student: [inaudible]

Professor Michael McBride: Ah, so when you run out of air in one, you switch the valve and take it from the other one; then you can go back to the first one. So you have an infinite supply of air, plenty to burn the oil. Right? But you need someone to control that, and you need someone else to make sure the lamp stays lit, that it gets the proper amount of oil siphoned in, and the airflow is right. You need someone else — in the CO2 collectors, you just blow gas over the surface. Right? So it’s got a little key on the top and a hook that goes down in. Somebody has to be standing there, doing this, to keep it stirred. Right? So that’s stirred. So it takes at least three or four people to run this thing. Right? And it cost a lot of money to build. So no one else built one. In fact, it’s not absolutely certain that Lavoisier ever used this, although it exists, right? Because he was busy with other things, having to do with the French Revolution, by the time this came along. But this was big science. Right? He was very wealthy, so he was able to do this kind of thing. But he showed, in this diagram, exactly how it could be done. So it was no secret, not like what the alchemists did. Was there a question here? No. Yes Lucas?

Student: Did they ever measure like the air supply they gave it?

Professor Michael McBride: Did they measure —

Student: Did they measure the amount of air that they gave into it? Did they know like the amount of the air?

Professor Michael McBride: I don’t think so, because there wasn’t any purpose; you didn’t care how much air you put in. Because most of the air just flowed right on through, it was never collected. Right? It just carried the other gases along, where the water and the CO2 would be collected.

Student: Well I guess if they knew the content of the air, like they could just measure how much oxygen-

Professor Michael McBride: To get the oxygen? Conceivably; but that was never done. It would’ve been very difficult to collect all that air and measure it accurately. Okay now what if you have a substance that won’t burn cleanly, so you can’t measure these gases? Notice, incidentally, the role of gases in understanding things, because they provide easy separation. Right? And then since he was able to weigh gas, once, then he knows the gas density. So now all you have to do, from then on, in principle, is measure the volume of gas, and you know how much stuff there was. So gases played a really key role through all this.

But suppose you had a substance that wouldn’t burn, that you couldn’t convert to steam and CO2. What do you do then? For example, grape sugar. So here’s chapter thirteen where he talks about le suc des raisins, the grape sugar; which will char, it won’t burn cleanly. So what do you do? Well, as he says up at the top there, “tout le monde sait comment se fait le vin”; everyone knows how to make wine, cider and mead. Right? So can you see what he’s going to do? So here’s the device for that one, which is in many ways sort of similar, but you don’t need an air — you don’t put the air in. What you have is a big flask at the beginning which has sugar, yeast and water. Right? So the yeast ferments the sugar and changes it into CO2. So you have — it foams up; if you’ve even seen something fermenting. So you have to have something that catches the foam so it won’t get into your device. And then you have something that will absorb water. Now the water is not just — it’s not like the previous one, where the water is all formed by the thing, because you got a whole bunch of water there at the beginning. Right? So that’s not what you’re trying to measure. What you’re trying to measure is how much carbon, how much CO2 there is in the grape sugar. So you have these sodium hydroxide solutions, the base that will absorb the CO2, as a carbonate. So you weigh those before and after. And the excess gas, if there’s another gas being evolved, will collect over the mercury in the far tube here. So he could measure the CO — the carbon, in sugar, by converting it completely to CO2 through fermentation.

Chapter 2. The Correct Experiment: Early Dealings with Experimental Data [00:12:25]

In speaking of fermentation here, he says it furnishes a means of analyzing sugar. So oxidation failed with air. Even with oxygen or sulfuric acid or mercuric oxide, you couldn’t get clean combustion of sugar; you got charring instead. But as he says here — this is the most interesting thing about his treatment of the data — he said, “I could consider the materials subjected to fermentation, and the products of fermentation, as an algebraic equation, having to do with weights.” Can you see what his assumption’s going to be here?

Student: [inaudible]

Professor Michael McBride: Sam?

Student: The same weights on both —

Professor Michael McBride: The same weights on both side; conservation of mass is his assumption. Okay? So he can make an equation that the weight on the right and the weight on the left will be the same. Okay, so an algebraic equation. “And by in turn, supposing each of the elements of the equation to be unknown, I can derive a value, and thus correct experiment by calculation, and calculation by experiment.” So what does he mean by that? So if you have this equation, a certain amount of sugar and a certain amount of oxygen gives a certain amount of CO2, for example. Right? So the weights on the two sides should be the same. He doesn’t know one of them. Right? But if he knows all the others, and it’s an algebraic equation, he can then solve for that one. So he says here: “I can derive a value and thus correct experiment by calculation.” Have you ever heard of that being done? Lexie?

Student: [inaudible]

Professor Michael McBride: Would somebody — would you do that in lab, correct experiment? Do you suppose there might be any nefarious people who would correct their experiments by calculation after finishing a lab?

Student: [laughs]

Professor Michael McBride: No, I can’t imagine that would happen, right? Okay, but Lavoisier was upfront. Right? He was even, one might say, bragging about it, that he has this algebraic equation, so he can correct his experiments. Right? “I’ve often profited from this way of correcting the preliminary results of my experiments.” No one told him that that’s not the way science works. Right? That the experiments are the things you have to do. You have to adjust the theory to deal with the experiments. Okay? Remember, he was the treasurer for the firm that collected taxes for the French Crown. Right? So he was accustomed to bookkeeping, and everything had to balance, right down to the sou. Right? So the same thing he applied to chemistry. So Table One here is the materials of fermentation. He put in a certain amount of water, 400 pounds; a certain amount of sugar, 100 pounds; yeast. Now he knew that the yeast was wet. So there was a certain amount of water. He put ten pounds of yeast, but he knew by analysis, by proximate analysis — right? — not what the elements are but what the substances are there — that in ten pounds of yeast, wet yeast, you’d have seven pounds, three ounces, six gros and forty-four grains. It turns that a gros is a 72nd of a grain, and eight gros make an ounce. So it’s a gros is — which is 28.35 grams. So, if you want, you can work through all this.

Do you think he had analyzed things that accurately, that he knew right down to the gros how much there was in there? No. There’s no way he could’ve known that. But things had to balance, right? So he made an approximate thing, and then he carried it out to eighteen decimal points, the way you’d do on a calculator. Right? Okay, so he didn’t know about significant figures yet either. Okay? So then in Table Two he figures in that much water, sugar, yeast, how much is there of — in each one, how much is in water, how much is hydrogen, and how much is oxygen. So he knows that. So now he knows down to hundredths of gros how much of each of those elements there are in the 400 pounds of water he’s using. Incidentally, he never used 400 pounds of water. He just rounded his numbers up, or multiplied his numbers, scaled them up to be that big. He used large quantities, but not that large. Okay.

So if you had 100 pounds of sugar, you’d have a certain amount of hydrogen, oxygen and charcoal - that is, carbon. And if you had a certain amount of the dry yeast you’d have that. And then he makes still a third table where he rearranges it, so he pulls the elements all together. So in the starting material, how much hydrogen, oxygen and carbon? Okay? So he’s got 411 pounds, 12 ounces, 6 gros, 1.36 grains of oxygen — grains and gros, got it; oh gros and grains, okay, right — okay, that much oxygen and some amount of hydrogen and charcoal, and a certain amount of nitrogen that’s in the yeast, because he knew that was there; which is just a small weight altogether, because there’s not much yeast. And in all, 510 pounds exactly. Right? So his balance sheet was just like what a banker would want. Okay, and now he does — he looks at the products, and he has a certain amount of CO2 that he collected; that’s carbonic acid. And he knows how much oxygen and charcoal; there is that. A certain amount of water. A certain amount of alcohol, if dry. So he had to collect the alcohol and say how much alcohol and how much water was in it. How much acetic acid. How much residual sugar that had fermented. And how much dry yeast there was at the end. And he added them all together, and by God, there was 510 pounds, zero ounces, zero gros, and zero grains. Right? QED. Okay? And then he recapitulated it, according to how much oxygen, how much charcoal and how much hydrogen. And once again it balances out absolutely precisely. Okay, so he recapitulates it that way. And this is how he would do some of these things.

Okay? So this is a description of an experiment, on page eighty-eight to ninety-two of the Traité Élémentaire. Okay, so he’s got a fire burning here, and a glass tube that runs through the middle of the charcoal that’s burning. Right? And he’s got a flask cemented to it, on the right, which is over a burner. And in the tube he puts twenty-eight grains of carbon — not pieces but a weight of twenty-two grains. We saw what a grain was last time. That turns out to be 1.38 grams of carbon he’s got inside the tube. Okay, and now do you think the carbon will evaporate, incidentally, if you put it in flame? Do you know how? How do you know it won’t evaporate?

Student: It shouldn’t.

Professor Michael McBride: Where do you get carbon? What did he call it?

Student: Charcoal.

Professor Michael McBride: He called it charcoal. You get charcoal by the flame, right? So with burning without enough oxygen. So obviously it’s not going to evaporate, it’s just going to sit there. Okay, so in flask A there he puts water. So there’s water in it. And then he lights the fire underneath and the water begins to boil, and distills through the tube and collects over here in H. Right? There it comes. Okay? But did you notice what happened in the process? Want me to back up and do it again? What happened?

Student: The charcoal went away.

Professor Michael McBride: The charcoal went away. What happened? Well let’s see what he found. So, the twenty-eight grains of carbon was gone. Right? And in the end he got the water that he had put in at the start, but not all the water. It was 85.7 grains short. So all the carbon and some of the water has gone. Where did it go? Well it went out the tube. So out the tube, he said, came 144 cubic inches — and because he knows the density, he knows how much weight that is — 100 grains, exactly, of carbonic gas, CO2; we would call it CO2 — and 380 cubic inches of flammable gas. So one of them would be collected by limewater, the carbonic gas, and the other one would go right on through. What do you suppose flammable gas is?

Students: Hydrogen.

Professor Michael McBride: Hydrogen. So the oxygen went to the carbon, and the hydrogen became elemental hydrogen, in this process. Okay? Now, look at this, this is what’s neat. So he got 100 grams — 100 grains, pardon me, of carbonic gas, and 13.7 grains of hydro-gen, the stuff that gives water when you burn it, according to him. Right? Which is right, which is correct. And what do you notice, quantitatively?

Student: They match.

Professor Michael McBride: They exactly match, right down to the 10ths of grains. Okay? Now, if you recalculate, by modern theory, how much gas you should get from twenty-eight grains of carbon doing this, you should get 157 cubic inches and 313 cubic inches, instead of 144 and 380; you should get 103 grains of carbonic gas and 9.4 grains of flammable gas. So his experiments weren’t perfect. Right? Which is not surprising, and there’s no shame in that. But what would we now say he did wrong?

Student: He adjusted his results.

Professor Michael McBride: He adjusted his results to fit his theory. Okay? But this was early days. He thought, he actually said, on page ninety-two: “I have thought it best to correct by calculation and to present the experiment in all its simplicity.” Now what’s wrong with presenting an experiment in all its simplicity? You lose anything that’s new, anything that you might have discovered in the experiment. Okay? But he was discovering a lot of stuff, so we’ll forgive him that. Okay? He was certainly a great scientist.

So what were his contributions? Clarity, above all. Right? This idea of facts and ideas and words having to go together and improve together. And of the facts he worked on, there were these — the apparatus I showed, and others as well. The important thing was quantitation, using numbers, which people hadn’t used before — not nearly as effectively at least — concentrating on making things balance, using mass as the criterion for how much stuff you have. But also, in the case of gases, volume, but measuring the density, so that you can relate volume to mass, which is the fundamental quantity.

He found some new substances — although that wasn’t much of a contribution on his part; other people, like Scheele before him, had discovered lots and lots of new pure substances; that wasn’t Lavoisier’s strong suit — and reactions. And then, with respect to ideas, the idea of what an element is; that it’s not some platonic ideal, but it only means as far as we’ve so far been able to go in analyzing things and breaking them down. And if we can’t break them down further, we’ll assume for the time-being that they are elements, but maybe that they won’t be in the long run. The idea of conservation of mass, of course, is key. Oxidation as the organizing principle. The idea of a radical, that an element is the base of something and that it can become an acid through oxidation. And the idea of salts, from mixing of elements. And with respect to words, the idea that names should be meaningful; in particular, the names for elements, oxidation state. So he developed these suffixes; -ous, -ic, -ide, -ite and -ate are all due to Lavoisier. And the idea of a salt composition, the way of naming them, that sodium hydrox-ide, for example, or sodium chlor-ide, or chlor-ate, or chlor-ite; which are different degrees of oxidation of the element.

But, like everyone who would follow, and everyone before, and everyone nowadays, he was plagued by lack of imagination. Right? That’s always the short suit of everyone who’s in science, or probably in anything. He said, at the end of this preface that we’ve talked about: “Chemistry’s present progress, however, is so rapid, and the facts, under modern doctrine, have assumed so happy an arrangement, that we have ground to hope, even in our own times, to see it approach near to the highest state of perfection of which it is susceptible.” Right? So people always think that; that there are a few things you see you have to do to fill in little chinks here and there, but otherwise we’ve got it now. Right? This is certainly what a lot of people tended to think about the human genome. Once you know what the human genome is, then all diseases are solved, everything else. Right? It’s never been true and it almost certainly is not true now. And it certainly wasn’t true in his time. He said, “even in our own times.” He wrote this in 1789, and in 1794 he was guillotined. [Laughter] So it certainly didn’t happen in his own time. The judge, Coffinhal, is reported to have said, although perhaps apocryphally, that “the Republic has no need of geniuses.” And, in fact, all his equipment was seized for The People, including his 80 pounds of mercury that he carried around in this thing, to measure gases. That madness lasted only a couple of years; maybe not even that long. And his widow recovered the equipment, and it currently is in the Musée des Arts et Métiers in Paris. I’ve never seen it, but a number of students from the class have sent me postcards when they’ve gone to see it. Okay, and Lagrange, the guy that helped him build the calorimeter [correction: it was Laplace who collaborated on the calorimeter], wrote: “It took them only an instant to make this head fall, but 100 years may not suffice to reproduce one like it.” So what we’re going to look at is what the heads in the next 100 years were able to accomplish.

Chapter 3. John Dalton’s Proportions and Atomic Theory [00:28:05]

So today we’re going to do the first twenty, or twenty-five, or thirty years after Lavoisier, to see where things went. So we’ve looked at Boyle and his gas laws, and Lavoisier. And now we’re going to go on to Dalton. So here’s John Dalton, who lived up near the Lake District in England. And he was an amateur meteorologist. So he was interested in keeping weather records and so on. And in particular he was interested in why, in the atmosphere, if you collect air at the top of a mountain, and down in the bottom of a valley, you get the same composition. Why don’t the heavier gases sink and the lighter ones rise? Good question, right? Well people on the continent had proposed that different gases attract one another. So they tend to stick and therefore stay mixed. Right? But Dalton had a different idea. And this is a picture he drew of gases. The top is Azote, or nitrogen, and the bottom is hydrogen. So he said, “The atoms of one kind did not repel the atoms of another kind.” It’s not — well you’ll see in just a second.

Okay, so the center is his symbol for the atom; nitrogen in that case. And surrounding it, this sort of halo, is the “envelope” that’s heat associated with that particular atom. So these rays come out. And if the rays meet one another, between adjacent atoms in the gas, then they repel one another; according to Dalton’s theory. But if they don’t match, as in hydrogen and nitrogen, then they don’t repel one another. Right? So atoms of the same element tend to stay apart, but atoms of the other element are completely free to be inside there. They knew that the pressure went down, as you went up; that was no question, right? But why don’t they stratify? It’s because the different elements don’t repel one another. So he substituted homorepulsion, that the same atoms repel one another, for what the Europeans used as heteroattraction. Okay? Obviously that turned out to be nonsense. But Dalton developed three principles, in terms of these atoms that he had started using for this purpose of something about the atmosphere. These principles have come to be known as Definite Proportions, Equivalent Proportions, and Multiple Proportions.

So Definite Proportions is that pure compounds always have the same weight ratio of their elements. So you don’t get a sample, of sugar here, that’s a certain percentage of carbon and hydrogen, and a slightly different percentage of carbon and hydrogen here. And he explains it by saying that these compounds are a given ratio of elements. So you have that same ratio all the time when you have a particular compound.

Now Equivalent Proportions is a pretty simple idea, but it looks complex. Okay, so suppose you have a certain weight of A reacts with a certain weight of B; okay, to form AB. Okay? And a certain amount of A reacts with a certain amount of C, to form AC. But if it’s the same amount of A, then in some sense that weight of B and that weight of C are equivalent to one another. Okay? Because they react with the same weight of A; so they’re equivalent. But now suppose you have another compound D that also reacts with b parts of B. A reacted with b parts of B. Right? What does that tell you, about equivalence?

[Students speak over one another]

Professor Michael McBride: That means that a parts of A is equivalent to d parts of D. Okay? So just suppose two grams of A reacted with three grams of B, and two grams of A reacted with eight grams of C. That means three grams of B is equal to eight grams of C, for purposes of combining. Okay? And suppose 1.5 grams of D also reacted with three grams of B. Then 1.5 grams of D is equal to one gram of A. Okay? Now what could you conclude? What’s the final line going to be? What do you predict from this, if you have these kinds of equivalence; an equivalence between B and C and an equivalence between A and B? Can you see? Then you predict that if you react D and C, what’s going to happen? What weight ratio will they be in?

Student: d parts of —

Professor Michael McBride: Okay, d parts of D will react with c parts of C. Right? Because the D is like the A in the beginning. The C is like the B in the beginning. So d parts of D should react with c parts of C. That’s what multiple [correction: Equivalent] Proportions is. I’ll show you an example. Okay? And finally Multiple Proportions. If the same two elements form several different compounds, the weight ratios are related by simple factors. And I’ll show you an example of that. So these are all things that Dalton adduced as support for the idea that things were made of atoms, fundamental units.

Now Definite Proportions. Does the same compound always have the same ratio of the elements in it? The French School was divided on this question. Berthollet said, “Non!” Because he said they’re metal alloys, or natural “organic” materials that vary in their analysis. Right? Okay? And Proust said, “Oui!” Right? And that defined what chemistry was; that chemists then began to deal only with things that obeyed this Law of Multiple [correction: Definite] Proportions. They didn’t deal in fundamental ways with alloys or with complicated big molecules that had really complicated formulas, or couldn’t be separated as mixtures and therefore got different analyses at different times. They assumed that anything that was pure had this — obeyed the Law of Definite Proportions. So it was really more of a definition, than a law.

Okay, now Multiple Proportions. So there were oxides of carbon; you remember Lavoisier, 1789, defined 44% carbon and 56% oxygen in CO2. And again there are problems for you to work together to see how he did some of these things. So that’s one of them, I believe. And then in 1801 carbonic acid, as it was called, was analyzed and found to be twenty-one grams of carbon to seventy-two grams of oxygen. Okay? Or in the case of nitrogen, there were several oxides that had these ratios, all as analyzed in 1810. Right? Now, you can look at the weight ratios, if you want to. Oxygen to carbon ratio in the first one is 1.27, the other is 2.57. Or you can look in the oxides of nitrogen, and they’re 0.58, 1.27, 2.39. Right? Now, what do you notice about these? Do you learn anything from this? These guys got these results and sat down and scratched their heads. What did they come up with? Anything profound? Any relationships among these numbers, that you can see? Or are they just sort of random numbers? Yes. Nick?

Student: The ratio is kind of like —

Professor Michael McBride: Can you speak up a little bit?

Student: 1.27 is half of —

Professor Michael McBride: Ah, 1.27 is very close to half of 2.57. It’s not exact, but there’s going to be experimental error. Right? It’s about one to two; it’s actually 2.02. And if you look in the case of nitrogen, it’s 1:2:4. Right? It’s actually 2.19 and 4.12. So as always, when people develop things experimentally and observe regularities, there’s a question of what’s real and what’s just experimental error; and you have no way of knowing at the beginning. We have a way of knowing now because we know what the true weights are. So they’re integral values consistent with simple atomic ratios. So now we know what the percent errors were in these things. Okay, that the ratio of the error was about 5% or 4% off from what it should’ve been in the oxygen:nitrogen. One is minus 2%, the others are plus 11%. So within roughly 5% or so, they were doing a pretty job of the analysis. But they were able to see through to this Law of Multiple Proportions. And that’s what you would expect, if these things were made up of identical molecules with simple ratios of numbers of atoms in them. So this tended to support Dalton’s ideas of atoms.

Chapter 4. Berzelius’s Contributions to Modern Precise Analysis and the Atomic Weight Confusion [00:37:28]

Okay, so we’ve looked at Lavoisier. And now we’re on further into elemental analysis, doing accurate elemental analysis, and the idea of atoms and dualism. And we’ll see what that is right away. And to do that we’re going to focus on Berzelius, but also talk about Gay-Lussac. Okay, so there’s Faraday, whom we’ve talked about already; Berzelius and Gay-Lussac, and also Davy, in England. Okay, so here’s Berzelius in Sweden. Born in 1779. Notice on the previous slide, all those guys were born within a year, 1778, 1779. Another guy that was born at the same time didn’t make it to the list, but his statue is right below, who’s Benjamin Silliman; he was exactly the same time. Okay, so here’s Berzelius, who lived to be almost seventy years old, and did an awful lot of stuff. So he was very good at analysis, both of organic substances and of minerals. In a period of six years he accurately analyzed 2000 compounds for their elements. Okay? He developed really good atomic weights. Remember, we just saw that these previous ones were off by as much as 10 or 11%, but he got very good ones for fifty different elements. Okay? He studied electrolysis, which we will see, and used that to develop the theory of dualism and how — understanding reactions as double decompositions; and we’ll shortly see what that is. And he did a lot of teaching and writing. He was the most respected chemist worldwide. He chose in his picture — remember, when these people got these pictures painted, they always chose some emblem to put in, and he chose his textbook that he wrote in 1803 [correction: 1808]. But he was a very important editor and summarizer of the year’s progress in chemistry every year. And he developed the notation that we still use for composition. So you’ll see that.

Okay, notation for composition. We’re going to look at its evolution, from alchemy through Dalton to Berzelius, which is what we use, almost exactly. Okay, so in 1774, this is a table printed in Sweden of symbols for elements and compounds. And you see, for example, here’s vitriolum cupri, or vitriolum coeruleum. Do you know what about cupri; that’s the symbol there of copper, right? So why would it have been called coeruleum? Can you see any word that’s related to that?

Student: Cerulean.

Professor Michael McBride: What’s cerulean about it?

Student: The color.

Professor Michael McBride: I got some here. That’s not it. That’s not it. This is it. It’s cerulean, right? Like the sky. Okay. Or iron, there’s the alchemical symbol of iron. Or and vitriol means sulfate, right?; or mercury sulfate; okay, lead. Okay? So these are the chemical symbols. Notice particularly nitrum, at the top, that comes from nitrates. Okay, and we’ll focus on that and see that that’s exactly the same symbol that Dalton used for nitrogen. So he derived his symbols for atoms from the alchemical symbols; at least many of them. Not all of them. Here are some others. And here are the symbols that Berzelius used for the same elements, which you see are the ones we use now. Okay? So notice as — so there were alchemical symbols at the beginning for Dalton, as you went across, but in the next row, once you get to iron, which is Fe in Berzelius, can you see what it is? It’s a little hard to read at this resolution. It’s a little circle with a letter inside. What’s the letter for iron? Can you see it? It’s in the middle of the second row. It’s an ‘I’. I’m sorry it’s not such a good resolution. And the next one is Z and then C and the L. Why? Z?

Student: Zinc.

Professor Michael McBride: Zinc. C?

Student: Copper.

Professor Michael McBride: Copper. L?

Student: Lead.

Professor Michael McBride: Lead. And so on. Okay, now down below that are binary symbols. So you put two elements together, like H with O, H with N, N with O, H with C, O with C. Now here’s Dalton’s logic about these symbols. He says, “When only one combination of two bodies can be obtained, it must be presumed to be a binary one.” That is, if you get only one compound, it must be one atom of one and one atom of the other. Right? So what would water be?

Student: H.

Professor Michael McBride: HO, or his circle with a dot in the circle. Right? “Unless some other cause appear to the contrary” that would show you that that’s not what you should assume. Okay, the next row, the ternary ones: N2O, NO2, CO2, CH2. And what he said about that is: “When two combinations are observed, of the same elements, then they must be presumed to be a binary and a ternary.” So you could have NO and NO2. What if there are three combinations of the same elements that give compounds? Do you see what he’s going to observe, what he’s going to suppose?

Student: Binary.

Professor Michael McBride: Right? “When three… a binary, and the other two ternary.” Okay? Okay, and what about if there are four? “When four… one binary, two ternary, and one quaternary.” Okay? So he’s laying down rules of logic for what, the way atoms must behave. And it turns out, as we know, atoms have different ideas from Dalton’s logic. There’s nothing illogical about Dalton, it’s just wrong. Okay, and this is Berzelius’s notation for the same symbols that Dalton gave on the left. But you notice they’re completely recognizable to them, except for one thing. What’s different?

Student: Superscripts instead of subscripts.

Professor Michael McBride: Here was ours, and here is Berzelius’s. The only difference is he used superscripts instead of subscripts. Okay? And he decided to base the names on Latin, so it would be international, whereas Dalton had used English. And the formulas, notice, are not structural. Berzelius’s aren’t, they’re just written in a line and the number. These could’ve been assumed to be structural, that you thought the atoms were arranged that way. I don’t know that Dalton really had that idea, but they sort of convey that idea. But Berzelius’s don’t, they’re just a list of the elements and the ratio, by weight, of atoms in them. And he developed some more fancy abbreviations as well, where, for example, these little dots above elements would who show how much oxygen was associated with it. So that chromate potassique was K with an O, and chromium with three Os. Or down here he has some even fancier things. Notice the superscripts denote the number of atoms. And this benzoic acid should be H10C14O3. Now wait a second, a carboxylic acid has how many oxygens? Carboxylate group?

Students: Two.

Professor Michael McBride: Two. Here he’s got three. The reason is that he felt there was water in these acids. So he heated the heck out of them to drive the water out, and what you actually were left with was the anhydride. So two acids had come together and lost water. Right? So the formula should be C[H]10C14O3, and he’s got H12C15. So not so far off. And he’s got these neat symbols for acetic, tartaric, citric and so on, acid. But those didn’t catch on either. But his symbols for the elements did.

Now let’s see what we got here, I had one more topic. Atomic weights and equivalents. Okay, these are Dalton’s atomic weights that he used in 1808: hydrogen one, carbon five, nitrogen five, oxygen seven, and so on. Now as of 2004, this is what they should’ve been. What’s wrong with Dalton? Why is he off so far, for example, in oxygen, which should be sixteen, and he has seven? How can he be off that much? Charles? Can’t hear.

Student: Diatomic.

Professor Michael McBride: Yeah, what did he think water was?

Student: HO.

Professor Michael McBride: HO. We know it’s H2O. So the ratio is off by a factor of two. Right? So all these are off — most of these are off by a factor of two or three, depending on what the true formula we now know to be. Okay? But if you make that correction, this is what kind of errors he had in his atomic weights. There were roughly 10% or so errors. Right? And the question in silicon chloride. Thomas Thomson in Scotland thought it was SiCl. Gmelin in Germany thought it was SiCl2. Berzelius thought it was SiCl3 and Odling in England thought it was SiCl4. So you could have anything you wanted. Right? So which one is it? How can you tell which one it is? That’s the problem. Okay, this is when Gay-Lussac in France gets into the act. So in 1804, he and the physicist Biot made a balloon ascent to a height of 7000 meters. That’s way the heck up there. And in fact, that was a record for fifty years. And they have a big flask they’re carrying. Why? What do they want to do with their big flask?

Student: The flask was to collect air.

Professor Michael McBride: Collect air at 7000 meters and see if it’s different. Okay? Okay, so they established that the atmospheric composition is invariant with altitude; at least at the heights they could do. But that’s not the most important thing he did, for our purpose. One of the really important things he did was figure out how to oxidize sugar, how to burn sugar cleanly. Remember, Lavoisier had to resort to fermentation. But he found if the source of the oxidation was sodium chlorate, then you could heat them together and actually get it right and get water and CO2 out of sugar. So that was a very important practical contribution. For the present purposes, he found that when you decompose water you get, to one volume of oxygen you get 1.9989 volumes of hydrogen. And if you do the same thing with ammonia, for a volume of nitrogen you get 3.08163 volumes of hydrogen, in his experiment. What did he conclude? What was Gay-Lussac’s conclusion for this?

Student: H2.

Professor Michael McBride: What do you think he thought the values would’ve been if there was no experimental error? How much hydrogen would he have gotten?

Student: Two.

Professor Michael McBride: Two. And how much hydrogen would he have gotten from nitrogen?

Student: Three.

Professor Michael McBride: Three. Right? They’re very, very close, just a few percent off. Right? And even closer in the case of water. So what did he think? He thought the number, that the volume of a gas was proportional to the number of atoms in it. Right? Whose name do you associate that with?

Student: Avogadro.

Professor Michael McBride: Avogadro. But both of them did it, independently, and Gay-Lussac is the one whose data I’m showing here. Okay, so this was an alternative to Dalton’s Law of Greatest Simplicity. Dalton said if there’s only one compound, it must be one to one. Gay-Lussac said look at the gas ratios. And that turned out to be correct; although it didn’t get completely adopted for another sixty years. That’s really funny. Okay, that’s enough for now.

[end of transcript]

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