Equal Numbers in Equal Volumes: Avogadro Elements and Atoms: Chapter 9 Equal Numbers in Equal Volumes: Avogadro Amedeo Avogadro (1776-1856; see portrait at the Edgar Fahs Smith collection, University of Pennsylvania) was an Italian chemist and physicist best known for today for Avogadro's law and Avogadro's number. Avogadro did not discover or determine Avogadro's number; its determination occured late in the 19th century and early in the 20th and it was named in his honor. [ Perrin 1908 ] It is the number of atoms in 1 gram of hydrogen, or equivalently the number of atoms in one mole: 6.022x1023 . This chapter includes Avogadro's law, that equal volumes of gas contain equal numbers of particles, presented as a hypothesis. Combined with Dalton's atomic hypothesis ( chapter 7 ), the law explains Gay-Lussac's observations about combining volumes of gases ( chapter 8 ).

Avogadro's ideas removed an important obstacle to the determination of atomic weights by allowing molecular formulas to be determined independently of mass measurements, at least for volatile compounds. I should say that his ideas would have removed that obstacle if they were adopted by other chemists; however, they were not widely accepted for about 50 years, until Stanislao Cannizzaro made them the foundation of a logical and coherent system [ Cannizzaro 1858 ]. from Essay on a Manner of Determining the Relative Masses of the Elementary Molecules of Bodies, and the Proportions in Which They Enter into These Compounds Journal de Physique 73, 58-78 (1811), as translated in Alembic Club Reprint No. 4 , (Edinburgh, 1890) I. M. Gay-Lussac has shown in an interesting Memoir (Mémoires de la Société d'Arcueil, Tome II.) that gases always unite in a very simple proportion by volume, and that when the result of the union is a gas, its volume also is very simply related to those of its components.

But the quantitative proportions of substances in compounds seem only to depend on the relative number of molecules which combine, and on the number of composite molecules which result. [1] It must then be admitted that very simple relations also exist between the volumes of gaseous substances and the numbers of simple or compound molecules which form them. The first hypothesis to present itself in this connection, and apparently even the only admissible one, is the supposition that the number of integral molecules in any gases [sic] is always the same for equal volumes, or always proportional to the volumes. [2] Indeed, if we were to suppose that the number of molecules contained in a given volume were different for different gases, it would scarcely be possible to conceive that the law regulating the distance of molecules could give in all cases relations as simple as those which the facts just detailed compel us to acknowledge between the volume and the number of molecules.

[3] On the other hand, it is very well conceivable that the molecules of gases being at such a distance that their mutual attraction cannot be exercised, their varying attraction for caloric may be limited to condensing the atmosphere formed by this fluid having any greater extent in the one case than in the other, and, consequently, without the distance between the molecules varying; or, in other words, without the number of molecules contained in a given volume being different. Dalton, it is true, has proposed a hypothesis directly opposed to this, namely that the quantity of caloric is always the same for the molecules of all bodies whatsoever in the gaseous state, and that the greater or less attraction for caloric only results in producing a greater or less condensation of this quantity around the molecules, and thus varying the distance between the molecules themselves.

[3a] But in our present ignorance of the manner in which this attraction of the molecules for caloric is exerted, there is nothing to decide us à priori in favour of the one of these hypotheses rather than the other; and we should rather be inclined to adopt a neutral hypothesis, which would make the distance between the molecules and the quantities of caloric vary according to unknown laws, were it not that the hypothesis we have just proposed is based on that simplicity of relation between the volumes of gases on combination, which would appear to be otherwise inexplicable. [4] Setting out from this hypothesis, it is apparent that we have the means of determining very easily the relative masses of the molecules of substances obtainable in the gaseous state, and the relative number of these molecules in compounds; for the ratios of the masses of the molecules are then the same as those of the densities of the different gases at equal temperature and pressure, and the relative number of molecules in a compound is given at once by the ratio of the volumes of the gases that form it.

[5] For example, since the numbers 1.10359 and 0.07321 express the densities of the two gases oxygen and hydrogen compared to that of atmospheric air as unity, and the ratio of the two numbers consequently represents the ratio between the masses of equal volumes of these two gases, it will also represent on our hypothesis the ratio of the masses of their molecules. Thus the mass of the molecule of oxygen will be about 15 times that of the molecule of hydrogen, or, more exactly as 15.074 to 1. In the same way the mass of the molecule of nitrogen will be to that of hydrogen as 0.96913 to 0.07321, that is, as 13, or more exactly 13.238, to 1. On the other hand, since we know that the ratio of the volumes of hydrogen and oxygen in the formation of water is 2 to 1, it follows that water results from the union of each molecule of oxygen with two molecules of hydrogen. [6] Similarly, according to the proportions by volume established by M.

Gay-Lussac for the elements of ammonia, nitrous oxide, nitrous gas , and nitric acid , ammonia will result from the union of one molecule of nitrogen with three of hydrogen, nitrous oxide from one molecule of oxygen with two of nitrogen, nitrous gas from one molecule of nitrogen with one of oxygen, and nitric acid from one of nitrogen with two of oxygen. [7] II. There is a consideration which appears at first sight to be opposed to the admission of our hypothesis with respect to compound substances. It seems that a molecule composed of two or more elementary molecules should have its mass equal to the sum of the masses of these molecules; and that in particular, if in a compound one molecule of one substance unites with two or more molecules of another substance, the number of compound molecules should remain the same as the number of molecules of the first substance. Accordingly, on our hypothesis, when a gas combines with two or more times its volume of another gas, the resulting compound, if gaseous, must have a volume equal to that of the first of these gases. Now, in general, this is not actually the case. [8] For instance, the volume of water in the gaseous state is, as M.

Gay-Lussac has shown, twice as great as the volume of oxygen which enters into it, or, what comes to the same thing, equal to that of the hydrogen instead of being equal to that of the oxygen.

[9] But a means of explaining facts of this type in conformity with our hypothesis presents itself naturally enough; we suppose, namely, that the constituent molecules of any simple gas whatever (i.e., the molecules which are at such a distance from each other that they cannot exercise their mutual action) are not formed of a solitary elementary molecule, but are made up of a certain number of these molecules united by attraction to form a single one; [10] and further, that when molecules of another substance unite with the former to form a compound molecule, the integral molecule which should result splits itself into two or more parts (or integral molecules) composed of half, quarter, &c., the number of elementary molecules going to form the constituent molecule of the first substance, combined with half, quarter, &c., the number of constituent molecules of the second substance that ought to enter into combination with one constituent molecule of the first substance (or, what comes to the same thing, combined with a number equal to this last of half-molecules, quarter-molecules, &c., of the second substance); so that the number of integral molecules of the compound becomes double, quadruple, &c., what it would have been if there had been no splitting up, and exactly what

 is necessary to satisfy the volume of the resulting gas.

[11] Thus, for example, the integral molecule of water will be composed of a half-molecule of oxygen with one molecule, or what is the same thing, two half-molecules of hydrogen. On reviewing the various compound gases most generally known, I only find examples of duplication of the volume relatively to the volume of that one of the constituents which combines with one or more volumes in the other. We have already seen this for water. In the same way, we know that the volume of ammonia gas is twice that of the nitrogen which enters into it. M. Gay-Lussac has also shown that the volume of nitrous oxide is equal to that of the nitrogen which forms part of it, and consequently is twice that of the oxygen. Finally, nitrous gas, which contains equal volumes of nitrogen and oxygen, has a volume equal to the sum of the two constituent gases, that is to say, double that of each of them. [12] Thus in all these cases there must be a division of the molecule into two; but it is possible that in other cases the division might be into four, eight, &c.

The possibility of this division of compound molecules might have been conjectured à priori; for otherwise the integral molecules of bodies composed of several substances with a relatively large number of molecules, would come to have a mass excessive in comparison with the molecules of simple substances. We might therefore imagine that nature had some means of bringing them back to the order of the latter, and the facts have pointed out to us the existence of such means. [13] Besides, there is another consideration which would seem to make us admit in some cases the division in question; for how could one otherwise conceive a real combination between two gaseous substances uniting in equal volumes without condensation, such as takes place in the formation of nitrous gas? Supposing the molecules to remain at such a distance that the mutual attraction of those of each gas could not be exercised, we cannot imagine that a new attraction could take place between the molecules of one gas and those of the other. But on the hypothesis of division of the molecule, it is easy to see that the combination really reduces two different molecules to one, and that there would be contraction by the whole volume of one of the gases if each compound molecule did not split up into two molecules of the same nature. M.

Gay-Lussac clearly saw that, according to the facts, the diminution of volume on the combination of gases cannot represent the approximation of their elementary molecules. [14] The division of molecules on combination explains to us how these two things may be made independent of each other. III. Dalton, on arbitrary suppositions as to the most likely relative number of molecules in compounds , has endeavoured to fix ratios between the masses of the molecules of simple substances. [15] Our hypothesis, supposing it well-founded, puts us in a position to confirm or rectify his results from precise data, and, above all, to assign the magnitude of compound molecules according to the volumes of the gaseous compounds, which depend partly on the division of molecules entirely unexpected by this physicist. Thus Dalton supposes that water is formed by the union of hydrogen and oxygen, molecule to molecule. [16] From this, and from the ratio by weight of the two components, it would follow that the mass of the molecule of oxygen would be to that of hydrogen as 71/2 to 1 nearly, or, according to Dalton's evaluation, as 6 to 1.

This ratio on our hypothesis is, as we saw, twice as great, namely, as 15 to 1. As for the molecule of water, its mass ought to be roughly expressed by 15+2=17 (taking for unity that of hydrogen), if there were no division of the molecule into two; but on account of this division it is reduced to half, 81/2, or more exactly 8.537, as may also be found by dividing the density of aqueous vapour 0.625 (Gay-Lussac) by the density of hydrogen 0.0732. [17] This mass differs from 7, that assigned to it by Dalton, by the difference in the values for the composition of water; so that in this respect Dalton's result is approximately correct from the combination of two compensating errors,--the error in the mass of the molecule of oxygen, and his neglect of the division of the molecule. [18] Dalton supposes that in nitrous gas the combination of nitrogen and oxygen is molecule to molecule: we have seen on our hypothesis that this is actually the same.

Thus Dalton would have found the same molecular mass for nitrogen as we have, always supposing that of hydrogen to be unity, if he had not set out from a different value for that of oxygen, and if he had taken precisely the same value for the quantities of the elements in nitrous gas by weight. But supposing the molecule of oxygen to be less than half what we find, he has been obliged to make that of nitrogen also equal to less than half the value we have assigned to it, viz., 5 instead of 13. [19] As regards the molecule of nitrous gas itself, his neglect of the division of the molecule again makes his result approach ours; he has made it 6+5=11, whilst according to us it is about , or more exactly , as we also find by dividing 1.03636, the density of nitrous gas according to Gay-Lussac, by 0.07321. Dalton has likewise fixed in the same manner as the facts has given us, the relative number of molecules in nitrous oxide and in nitric acid, and in the first case the same circumstance has rectified his result for the magnitude of the molecule.

He makes it 6 + 2x5 = 16, whilst according to our method it should be , a number which is also obtained by dividing 1.52092, Gay-Lussac's value for the density of nitrous oxide, by the density of hydrogen. In the case of ammonia, Dalton's supposition as to the relative number of molecules in its composition is on our hypothesis entirely at fault. He supposes nitrogen and hydrogen to be united in it molecule to molecule, whereas we have seen that one molecule of nitrogen unites with three molecules of hydrogen. According to him the molecule of ammonia would be 5+1=6; according to us it should be , or more exactly 8.119, as may also be deduced directly from the density of ammonia gas. The division of the molecule, which does not enter into Dalton's calculations, partly corrects in this case also the error which would result from his other suppositions. All the compounds we have just discussed are produced by the union of one molecule of one of the components with one or more molecules of the other.

In nitrous acid we have another compound of two of the substances already spoken of, in which the terms of the ratio between the number of molecules both differ from unity. From Gay-Lussac's experiments (Société d'Arcueil, same volume) it appears that this acid is formed from 1 part by volume of oxygen and 3 of nitrous gas [20] , or, what comes to the same thing, of 3 parts of nitrogen and 5 of oxygen; hence it would follow, on our hypothesis, that its molecule should be composed of 3 molecules of nitrogen and 5 of oxygen, leaving the possibility of division out of account. But this mode of combination can be referred to the preceding simpler forms by considering it as the result of the union of 1 molecule of oxygen with 3 of nitrous gas, i.e. with 3 molecules, each composed of a half-molecule of oxygen and a half-molecule of nitrogen, which thus already included the division of some of the molecules of oxygen which enter into that of nitrous acid.

Supposing there to be no other division, the mass of this last molecule would be 57.542, that of hydrogen being taken as unity, and the density of nitrous acid gas would be 4.21267, the density of air being taken as unity. But it is probable that there is at least another division into two, and consequently a reduction of the density to half: we must wait until this density has been determined by experiment. [21] ... [22] VIII. It will have been in general remarked on reading this Memoir that there are many points of agreement between our special results and those of Dalton, although we set out from a general principle, and Dalton has only been guided by considerations of detail. [23] This agreement is an argument in favour of our hypothesis, which is at bottom merely Dalton's system furnished with a new means of precision from the connection we have found between it and the general fact established by M. Gay-Lussac. Dalton's system supposes that compounds are made in general in fixed proportions, and this is what experiment shows with regard to the more stable compounds and those most interesting to the chemist.

It would appear that it is only combinations of this sort that can take place amongst gases, on account of the enormous size of the molecules which would result from ratios expressed by larger numbers, in spite of the division of the molecules, which is in all probability confined within narrow limits. We perceive that the close packing of the molecules in solids and liquids, which only leaves between the integral molecules distances of the same order as those between the elementary molecules, can give rise to more complicated ratios, and even to combinations in all proportions; but these compounds will be so to speak of a different type from those with which we have been concerned, and this distinction may serve to reconcile M. Berthollet's ideas as to compounds with the theory of fixed proportions. [24] Notes [1] Avogadro uses the term molecule and never atom. His molecules are similar conceptually to Dalton's atoms, the smallest piece of matter which retained its identity as a particular substance ( chapter 7, note 5 ). Sometimes he modifies the word, as in composite molecule or (below) elementary molecule. An "elementary molecule" is what a modern chemist would call an atom.

A "composite molecule" is what a modern chemist would call a molecule (i.e., two or more atoms bound together). [2] This is Avogadro's law, presented as a hypothesis. In modern terms, it states that equal volumes of gas contain equal numbers of molecules. [3] Why might one think that equal volumes of gas contain equal numbers of molecules? For one thing, it is the simplest explanation of Gay-Lussac's observations that gases combine in simple ratios by volume. For if atoms combine in simple ratios by volume (as Dalton believed), and if equal volumes of gas contain equal numbers of molecules, then it would follow that gases would combine in simple ratios by volume.The close relationship between Gay-Lussac's observations and Dalton's ideas was not sufficient, however, to convince Dalton either of Gay-Lussac's observations or of Avogadro's hypothesis.

(See chapter 8, note 42 .) In fact, the hypothesis of equal volumes containing equal numbers of molecules had already occured to Dalton, but he rejected it: "At the time I formed the theory of mixed gases, I had a confused idea, as many have, I suppose, at this time, that the particles of elastic fluids are all of the same size; that a given volume of oxygenous gas contains just as many particles as the same volume of hydrogenous; ..." [Dalton 1808, p. 188]. The reasons Dalton rejected the hypothesis had to do with caloric supposed to be associated with gaseous particles. [3a] Note that Avogadro has already referred to the work of two of his contemporaries. Scientists typically build upon, criticise, and modify the work of their predecessors and contemporaries. In this sense, they do not work alone. Often the interaction among scientists and their ideas is international. Here the Italian (or, more accurately, Piedmontese, as Italy was not yet unified) Avogadro refers to work of the English Dalton and the French Gay-Lussac.

[4] Avogadro accepts the conventional wisdom of his time that molecules in a gas are coated with caloric (heat), which overcomes the forces of cohesion apparent in solids and liquids. Dalton pictures this caloric [Dalton 1808] as a mechanical spacer keeping molecules apart, the distance of separation determined by a substance's affinity to caloric. Avogadro maintains that not enough is known about caloric to decide the question; therefore, Dalton's picture should not be an obstacle to Avogadro's hypothesis. [5] Avogadro recognized that if his hypothesis was correct, it provided the basis for determining the relative atomic and molecular weights of gases and the formulas of gaseous compounds. If equal volumes of gas contain equal numbers of molecules, then their densities (mass per unit volume) are proportional to their molecular weights (mass per molecule). Ratios of gas densities, then, are equal to ratios of molecular weights. "The relative number of molecules [atoms] in a compound" is nothing more than the compound's formula, a list of the atoms in the compound and their relative numbers. Already Avogadro has begun to give some evidence for his hypothesis (previous paragraph) and to sketch out its implications.

[6] That is, the formula for water is H2O; compare to Dalton's assumed formula, HO. [7] That is, the formula for ammonia is NH3; compare to Dalton's assumed formula, NH. "Nitrous gas" is known today as nitric oxide, NO; "nitric acid" is known today as nitrogen dioxide, NO2. Dalton got right the formulas for the three compounds of nitrogen with oxygen. [8] Gay-Lussac presented some examples of these apparent paradoxes. See notes 39 and 40 of the previous chapter. [9] For example, Gay-Lussac reported that one part oxygen reacts with two parts hydrogen to give two parts water by volume. Why two parts of water, Avogadro asks: after all, if water is H2O, shouldn't the two particles of hydrogen attach to the one particle of oxygen to make one particle of water? [10] In modern terminology, Avogadro supposes that any gaseous element ("simple gas") exists not as single atoms ("solitary elementary molecule") but as molecules composed of more than one atom. This hypothesis turned out to be correct, but it was not widely accepted at the time.

Indeed, there were some good reasons for skepticism. Some scientists were not convinced of the atomic hypothesis in the first place [Brock & Knight 1967, Knight 1967]; they naturally were not convinced that gaseous elements existed as pairs of atoms bound together. In addition, many scientists who thought the atomic hypothesis did make sense (including Dalton) saw the idea of diatomic gases as nonsense. These scientists believed (correctly) that gases were made up of widely separated particles. It did not make sense that these particles would form pairs and that these pairs would then stay apart from other pairs: if atoms of the same kind were attracted to each other, they would all cluster together to form a solid or liquid; if they repelled each other (as they were believed to do when they were attached to caloric), then no pairs would form. Thus, the hypothesis of composite molecules for gaseous elements was one which had no direct evidence in its favor, while it had a considerable quantity of scientific opinion against it. The only point in its favor at the time was that it (in conjunction with the hypothesis of equal volumes of gas containing equal numbers of molecules) explained the available data.

Thus the notion of composite molecules was a classic case of an ad hoc hypothesis, one which was not intimately connected with the rest of the theory in which it appeared and which was introduced only to make the rest of the theory work. [11] Many commentators have remarked that Avogadro's writing in this paper is not always clear. Leonard Nash, for example, writes [Nash 1957], "The French text suffers from occasional obscurity, which must have confused Avogadro's contemporaries even as it does us." Perhaps the obscurity arises (as Nash hints) because Avogadro was not writing in his native Italian; perhaps because the ideas were not entirely clear in his mind. Perhaps the lack of clarity in Avogadro's prose contributed to the difficulty his ideas experienced in becoming established. What Avogadro is suggesting here may best be illustrated by example. Suppose, he says, that hydrogen and oxygen gases consist not of a single atom but of two atoms bound together. Using modern notation, we would say that these elements have formulas of H2 and O2 rather than H and O respectively.

Now suppose that when they react in the observed proportions of two parts hydrogen to one part oxygen, they form a larger molecule which then splits up to give two molecules of water as the ultimate products: 2 H2 + O2 --> [H4O2] --> 2 H2O . Today, we recognize Avogadro's two suppositions or hypotheses as correct: equal volumes of gas contain equal numbers of molecules, and many gaseous elements exist as diatomic molecules. In fact, by the end of the 19th century, when argon was discovered, the idea of a monatomic gas was shocking. (See chapter 14, note 52 .) The actual mechanism of water formation does not involve the formation and then splitting of a double molecule of water, H4O2, though. Avogadro recognizes that he cannot tell if the original molecules of hydrogen and oxygen contained two atoms or four atoms or some other even number of atoms.

For example, he recognized that if hydrogen and oxygen had the formulas H4 and O4, the reaction 2 H4 + O4 --> [H8O4] --> 2 H4O2 would still explain the data: hydrogen and oxygen would still combine in a ratio of 2 to 1; water would still have the ratio of two hydrogen to one oxygen; and the volume of water produced would be the same as the volume of hydrogen, not oxygen. Indeed, André-Marie Ampère proposed a hypothesis quite similar to Avogadro's except that it explicitly proposes tetratomic (4-atom) molecules [Ampère 1814]. In the absence of any additional information, however, the notion of diatomic molecules is more attractive than that of tetratomic molecules because it is simpler. [12] The reactions mentioned here are: N2 + 3 H2 --> 2 NH3 , 2 N2 + O2 --> 2 N2O , N2 + O2 --> 2 NO . [13] Avogadro seems to suggest that there is some sort of natural tendency to keep molecules from getting very large or complicated.

He would be surprised, then, by the knowledge chemists have built up about macromolecules (molecules consisting of hundreds or thousands of atoms). Both natural and artificial macromolecules abound. DNA (which carries the genetic code of living things) and hemoglobin (a protein which transports oxygen in the blood) are two examples of natural macromolecules of tremendous biological importance. Most modern plastics are examples of synthetic macromolecules. [14] That is, the diminution of volume in these reactions is much less than can be explained if the only thing going on was the atoms getting close together ("approximation") by attraction. [15] See chapter 7, note 18 . [16] Here and below, Avogadro means one molecule of each reactant by "molecule to molecule." [17] Recall that Avogadro proposed that a ratio of gas densities would be equal to a ratio of molecular weights. Here and in the remainder of this section, he compares the molecular weight of water based on gas densities to that based on his improved formulas. This is an attempt to compare different ways of measuring the same quantity.

In fact, the two ways of getting at the molecular weight of water are not quite independent; still, comparing the results shows that Avogadro's methods are at least self-consistent. [18] To summarize Avogadro's conclusions on water and its constituents: take the mass of H2 to be 1; then that of O2 is 15, and that of H2O is 81/2. Contrast this to Dalton's conclusions, which take the mass of H to be 1, that of O as 7, and that of HO as 8. A modern chemist would say that if the mass of H is taken to be 1, the mass of O is 16 and that of H2O is 18. The differences among these three sets of values stem from several sources. First, all use different analyses for the ratio of combining weights of hydrogen and oxygen in water (1:7 for Dalton, 1:7.5 for Avogadro; 1:8 today). These numbers represent the results of experiments which became progressively more accurate. They do not depend on the atomic hypothesis or Avogadro's hypotheses.

Avogadro's figures differ from Dalton's, besides the difference due to accuracy of the analysis, because Avogadro recognizes that water has two atoms of hydrogen for every atom of oxygen. Finally, Avogadro's numbers differ from modern numbers, besides the difference due to the accuracy of the analysis, because Avogadro in essence chooses H2 (the hydrogen molecule) as unity, while the modern system's basis is tantamount to taking the hydrogen atom as unity. [19] Actually, Dalton's assignment of 5 for nitrogen appears to be based on an analysis of ammonia, not of nitrous gas (NO). See chapter 7, note 20 for the uneasy coexistence of data on ammonia and on NO in Dalton's work. [20] This "nitrous acid" must be nitrous anhydride, N2O3. The reaction actually involves four times as much "nitrous gas" as oxygen, not three as stated: 4 NO + O2 --> 2 N2O3 . Nitrogen forms several compounds with oxygen, including NO, N2O3, NO2, N2O4, and N2O5, and frequently these gases appear in mixtures.

Therefore it is not surprising that Gay-Lussac's report was erroneous. [21] Avogadro suspects that the product of this reaction divides not just in two, as was the case in the examples he has already given, but in half again. Apparently Avogadro thought that such a large molecule was unlikely to be stable. (See footnote 13 above.) [22] I have omitted sections IV-VII of this paper. --CJG [23] Dalton, on the contrary, considered that his atomic hypothesis was general, and that conclusions drawn on the combining ratios of gases were restricted to gases. (Cf. chapter 8, note 40 .) I must agree with Dalton, with two provisos: Avogadro's ideas can be extended to volatile compounds which are not gases at room temperature; more important, Avogadro corrects a serious flaw in Dalton's system. [24] See previous chapter, note 6 regarding the question of definite proportions and variable proportions. Avogadro attempts to make a distinction between situations in which definite proportions apply and those in which variable proportions are permitted; however, it is not the modern distinction between compounds and mixtures ( chapter 8, note 40 ).

Reference Annales de chimie et de physique 90, 45- (1814); translated in Philosophical Magazine 45, 41-43, 109-116, 188-193, 344-349 (1815) , (Edinburgh, 1890); annotated excerpts above. W. H. Brock & D. M. Knight, "The Atomic Debates" in W. H. Brock, ed., The Atomic Debates (Leicester: Leicester U. Press, 1967), pp. 1-30. Stanislao Cannizzaro, 7, 321-366 (1858); translated as Alembic Club Reprints Avogadro's Law: Volume and Amount About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account.

There, you can easily access this resource later when you’re ready to customize it or assign it to your students. Avogadro's Law: Volume and Amount Avogadro's Law states that at the same temperature and pressure, equal volumes of different gases contain an equal number of particles. Learning Objective State Avogadro's Law and its underlying assumptions Key Points The number of molecules or atoms in a specific volume of ideal gas is independent of size or the gas' molar mass . Avogadro's Law is stated mathematically as follows: $\frac{V}{n} = k$ , where V is the volume of the gas, n is the number of moles of the gas, and k is a proportionality constant. Volume ratios must be related to the relative numbers of molecules that react; this relationship was crucial in establishing the formulas of simple molecules at a time when the distinction between atoms and molecules was not clearly understood. Term Full Text Definition of Avogadro's Law Avogadro's Law (sometimes referred to as Avogadro's hypothesis or Avogadro's principle) is a gas law; it states that under the same pressure and temperature conditions, equal volumes of all gases contain the same number of molecules.

The law is named after Amedeo Avogadro who, in 1811, hypothesized that two given samples of an ideal gas—of the same volume and at the same temperature and pressure—contain the same number of molecules; thus, the number of molecules or atoms in a specific volume of ideal gas is independent of their size or the molar mass of the gas. For example, 1.00 L of N2 gas and 1.00 L of Cl2 gas contain the same number of molecules at Standard Temperature and Pressure (STP). Avogadro's Law is stated mathematically as: $\frac{V}{n} = k$ V is the volume of the gas, n is the number of moles of the gas, and k is a proportionality constant. As an example, equal volumes of molecular hydrogen and nitrogen contain the same number of molecules and observe ideal gas behavior when they are at the same temperature and pressure. In practice, real gases show small deviations from the ideal behavior and do not adhere to the law perfectly; the law is still a useful approximation for scientists, however. The model contains gas molecules under constant pressure. The barrier moves when the volume of gas expands or contracts.

Run the model and select different numbers of molecules from the drop-down menu. What is the relationship between the number of molecules and the volume of a gas? (Note: Although the atoms in this model are in a flat plane, volume is calculated using 0.1 nm as the depth of the container.) Significance of Avogadro's Law Discovering that the volume of a gas was directly proportional to the number of particles it contained was crucial in establishing the formulas for simple molecules at a time (around 1811) when the distinction between atoms and molecules was not clearly understood. In particular, the existence of diatomic molecules of elements such as H2, O2, and Cl2 was not recognized until the results of experiments involving gas volumes was interpreted. Early chemists calculated the molecular weight of oxygen using the incorrect formula HO for water. This lead to the molecular weight of oxygen being miscalculated as 8, rather than 16. However, when chemists found that an assumed reaction of H + Cl $\rightarrow$ HCl yielded twice the volume of HCl, they realized hydrogen and chlorine were diatomic molecules.

The chemists revised their reaction equation to be H2 + Cl2 $\rightarrow$ When chemists revisited their water experiment and their hypothesis that $HO \rightarrow H + O$ , they discovered that the volume of hydrogen gas consumed was twice that of oxygen. By Avogadro's Law, this meant that hydrogen and oxygen were combining in a 2:1 ratio. This discovery led to the correct molecular formula for water (H2O) and the correct reaction $2H_2O \rightarrow 2H_2 + O_2$ Avogadro's Law - Avogadro's theory - Avogadro's hypothesis Avogadros law - What is it? Avogadro's Law (Avogadro's theory; Avogadro's hypothesis) is a principle stated in 1811 by the Italian chemist Amedeo Avogadro (1776-1856) that "equal volumes of gases at the same temperature and pressure contain the same number of molecules regardless of their chemical nature and physical properties". This number (Avogadro's number) is 6.022 X 1023.

It is the number of molecules of any gas present in a volume of 22.41 L and is the same for the lightest gas (hydrogen) as for a heavy gas such as carbon dioxide or bromine. The law can be stated mathematically . p is the pressure of the gas T is the temperature of the gas has the same value for all gases, independent of the size or mass of the gas molecules. One mole of an ideal gas occupies 22.4 liters (dm�) at STP, and occupies 24.45 litres at SATP (Standard Ambient Temperature and Pressure = 273K and 1 atm or 101.325 kPa). This volume is often referred to as the molar volume of an ideal gas. Real gases may deviate from this value. Or to put it another way "the principle that equal volumes of all gases at the same temperature and pressure contain the same number of molecules. Thus, the molar volume of all ideal gases at 0� C and a pressure of 1 atm. is 22.4 liters" Avogadro's number is one of the fundamental constants of chemistry.

It permits calculation of the amount of pure substance ( mole ), the basis of stoichiometric relationships. It also makes possible determination of how much heavier a simple molecule of one gas is than that of another, as a result the relative molecular weights of gases can be ascertained by comparing the weights of equal volumes. Avogadro's number (conventionally represented by N' in chemical calculations) is now considered to be the number of atoms present in 12 grams of the carbon-12 isotope (one mole of carbon 12) and can be applied to any type of chemical entity. Custom Search Amedeo Avogadro | Chemical Heritage Foundation Search Amedeo Avogadro Avogadro correctly hypothesized that equal volumes of gases, at the same temperature and pressure, contain equal numbers of molecules. Home / Learn / Historical Biographies In 1811 Avogadro put forward a hypothesis that was neglected by his contemporaries for years. Eventually proven correct, this hypothesis became known as Avogadro’s law, a fundamental law of gases.

The contributions of the Italian chemist Amedeo Avogadro (1776–1856) relate to the work of two of his contemporaries, Joseph Louis Gay-Lussac and John Dalton . Gay-Lussac’s law of combining volumes (1808) stated that when two gases react, the volumes of the reactants and products—if gases—are in whole number ratios. This law tended to support Dalton’s atomic theory, but Dalton rejected Gay-Lussac’s work. Avogadro, however, saw it as the key to a better understanding of molecular constituency. Avogadro’s Hypothesis In 1811 Avogadro hypothesized that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. From this hypothesis it followed that relative molecular weights of any two gases are the same as the ratio of the densities of the two gases under the same conditions of temperature and pressure. Avogadro also astutely reasoned that simple gases were not formed of solitary atoms but were instead compound molecules of two or more atoms. (Avogadro did not actually use the word atom; at the time the words atom and molecule were used almost interchangeably.

He talked about three kinds of “molecules,” including an “elementary molecule”—what we would call an atom.) Thus Avogadro was able to overcome the difficulty that Dalton and others had encountered when Gay-Lussac reported that above 100°C the volume of water vapor was twice the volume of the oxygen used to form it. According to Avogadro, the molecule of oxygen had split into two atoms in the course of forming water vapor. bio-avogadro.jpg Amedeo Avogadro. Courtesy of Edgar Fahs Smith Collection, Kislak Center for Special Collections, Rare Books and Manuscripts, University of Pennsylvania. Curiously, Avogadro’s hypothesis was neglected for half a century after it was first published. Many reasons for this neglect have been cited, including some theoretical problems, such as Jöns Jakob Berzelius ’s “dualism,” which asserted that compounds are held together by the attraction of positive and negative electrical charges, making it inconceivable that a molecule composed of two electrically similar atoms—as in oxygen—could exist.

In addition, Avogadro was not part of an active community of chemists: the Italy of his day was far from the centers of chemistry in France, Germany, England, and Sweden, where Berzelius was based. Personal Life Avogadro was a native of Turin, where his father, Count Filippo Avogadro, was a lawyer and government leader in the Piedmont (Italy was then still divided into independent countries). Avogadro succeeded to his father’s title, earned degrees in law, and began to practice as an ecclesiastical lawyer. After obtaining his formal degrees, he took private lessons in mathematics and sciences, including chemistry. For much of his career as a chemist he held the chair of physical chemistry at the University of Turin. The information contained in this biography was last updated on July 29, 2015. Ch 19 Unit 19 Gases and The Mole The purpose of this is to give quick reference to information or to use in an emergency (like if your text has accidentally been left under your desk at school). This is NOT intended to replace reading the text with its excellent photographs, diagrams, charts, and tables.

GASES AND THE MOLE In Chapter 18, we examined the effect of temperature and pressure on the volume of a constant mass of gas. The principles hold true for any gas exhibiting ideal behavior. Different gases, however, have molecules and atoms of different masses. In this chapter, we will look at the effect of the number of particles on the other gas variables, particularly volume. We will also find out how to obtain mass measurements of gases at various temperatures and pressures. 19:1 DEVELOPlNG AVOGADRO'S PRlNClPLE Suppose we place two different gases at exactly the same temperature and pressure in separate containers which have exactly the same volume. At a given temperature, all gas molecules will have the same average kinetic energy regardless of size or mass. Massive molecules will travel slowly, lighter molecules will travel more rapidly. However, the average kinetic energy (1/2 mv2) will be the same for all. If the kinetic energies are equal, any difference in pressure exerted by the gases is determined by the number of molecules of each gas. Since we have already said the two gases are at the same pressure, there must be an equal number of molecules in the two containers.

At equal temperatures and equal pressures, equal volumes of gases contain the same number of molecules. This statement is called Avogadro's principle, after Amadeo Avogadro. When he proposed the principle, in 1811, the kinetic theory had not been developed. As we read in Section 5:2, Avogadro developed his principle to explain some of the observations made by Gay-Lussac. Gay-Lussac had observed that two gases always react in such a way that the combining volumes can be expressed in small whole numbers. Proust had previously stated his law of definite proportions. Avogadro knew that one molecule of chlorine always united with one molecule of hydrogen. He also knew that two molecules of hydrogen chloride were formed. He therefore concluded that equal volumes of gases must contain equal numbers of molecules, Avogadro's principle has been verified so often that it is sometimes called a law. 19:2 MOLAR VOLUME One mole of any gas at STP will occupy the same volume as 1 mole of any other gas at STP.

This volume is 22.4 dm3, (22.4 liters/mol) the volume occupied by 1 mole of any gas under standard conditions, is called the molar volume of the gas. The molar volume for all gases at STP is 22.4 L/mol 19:3 lDEAL GAS EQUATlON We are now in a position to combine all four variables concerned with the physical characteristics of gases. These variables are pressure, volume, temperature, and number of particles. Charles' law states that volume varies directly as the absolute temperature. Boyle's law states that volume varies inversely as the pressure. Combining all the above factors gives us the ideal gas equation: PV = nRT The equation PV = nRT is called the ideal gas equation. The value of the new constant R can be obtained by substituting into the equation a set of known values of n, P, V, and T. We know that standard pressure P is 1O1 kPa, molar volume V is 22.4 dm3, standard temperature T is 273 K, and the number of moles n is 1.

So the constant R = 8.31 dm3kPa/molK This value for R is constant and can always be used if the units of the other quantities are not changed. We can determine the number of moles, n, in a quantity of a substance by dividing its mass by its formula mass. Any problem that can be solved by this equation can also be solved by direct application of the gas laws. SUMMARY 1. Avogadro's principle: At equal temperatures and pressures, equal volumes of gases contain equal numbers of molecules. 2. The molar volume of a gas is the volume occupied by 1 mole of the gas at STP; its value is 22.4 dm3 or 22.4 L/mol, for all gases. 3. The ideal gas equation is PV = nRT, where P = pressure, V = volume, n = number of moles, R = a constant, and T = kelvin temperature. 4. The molecular mass of a gas may be determined by using a modified form of the ideal gas equation. 5. A limiting reactant is a reactant which is completely consumed in a reaction. Any remaining reactants are said to be in excess. 6.

In solving problems involving gas volumes, consideration must be given to the change of gas volume with change in pressure and temperature. 7. A gas which expands or is compressed does work or has work done on it. More on Gases & Moles: Amedeo Avogadro - Biography, Facts and Pictures Amedeo Avogadro Lived 1776 – 1856. Amedeo Avogadro is best known for his hypothesis that equal volumes of different gases contain an equal number of molecules, provided they are at the same temperature and pressure. His hypothesis was rejected by other scientists. It only gained acceptance after his death. It is now called Avogadro’s law. He was also the first scientist to realize that elements could exist in the form of molecules rather than as individual atoms. Avogadro’s Life Amedeo Avogadro was born in Turin, Italy, on August 9th, 1776. His family background was aristocratic. His father, Filippo, was a magistrate and senator who had the title of Count. His mother was a noblewoman, Anna Vercellone of Biella. Amedeo Avogadro inherited the title of Count from his father.

In fact, Amedeo Avogadro’s full name was Count Lorenzo Romano Amedeo Carlo Avogadro di Quaregna e di Cerreto – quite a mouthful! Avogadro was highly intelligent. In 1796, when he was only 20, he was awarded a doctorate in canon law and began to practice as an ecclesiastical lawyer. Although he had followed the family tradition by studying law, he gradually lost interest in legal matters. He found science was much more intellectually stimulating. Mathematics and physics in particular attracted his logical mind. He spent increasing amounts of time studying these subjects. He was helped in this by the prominent mathematical physicist Professor Vassalli Eandi. Advertisements In 1803, in cooperation with his brother Felice, Avogadro published his first scientific paper, which looked at the electrical behavior of salt solutions. This was state-of-the-art science: only three years earlier, Avogadro’s fellow Italian scientist Alessandro Volta had demonstrated the electric battery for the first time. In 1806, aged 30, Avogadro abandoned his successful legal practice and started teaching mathematics and physics at a high school in Turin.

In 1809 he became a senior teacher at the College of Vercelli. In 1820 Avogadro became the professor of mathematical physics at the University of Turin. Unfortunately, this post was short lived because of political turmoil. Avogadro lost his job in 1823. Avogadro was reappointed in 1833 and remained in this post until, at the age of 74, he retired in 1850. Although he was an aristocrat, Avogadro was a down-to-earth, private man, who was quietly religious. He was dedicated to hard work and his lifestyle was simple. His wife’s name was Felicita Mazzé. They married in 1818 when Avogadro was aged 42. They had six sons. Avogadro’s Contributions to Science In the early 1800s, scientists’ ideas about the particles we now call atoms and molecules were very limited and often incorrect. Avogadro was deeply interested in finding out how the basic particles of matter behaved and came together to form chemical compounds. He studied the work of two other scientists: 1.

John Dalton In 1808 John Dalton published his atomic theory proposing that all matter is made of atoms. He further stated that all atoms of an element are identical, and the atoms of different elements have different masses. In doing so, Dalton carried chemistry to a new level. But he also made mistakes about the way elements combine to form compounds. For example, he thought water was made of one hydrogen atom and one oxygen atom and wrote it as HO; today we know water contains two hydrogens to every oxygen and we write water as H20. Actually, Avogadro figured this out, as we shall see. 2. Joseph Gay-Lussac In 1809 Joseph Gay-Lussac published his law of combining gas volumes. He had noticed that when two liters of hydrogen gas react with one liter of oxygen gas, they form two liters of gaseous water. All gases that he reacted seemed to react in simple volume ratios. Avogadro’s Hypothesis In 1811 Avogadro published a paper in Journal de Physique, the French Journal of Physics.

He said that the best explanation for Gay-Lussac’s observations of gas reactions was that equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules. This is now called Avogadro’s law. He published it when he was working as a physics teacher at the College of Vercelli. In Avogadro’s (correct) view, the reason that two liters of hydrogen gas react with a liter of oxygen gas to form just two liters of gaseous water is that the volume decreases because the number of particles present decreases. Therefore the chemical reaction must be: 2H2 (gas) + O2 (gas) → 2H20 (gas) In this reaction three particles (two hydrogen molecules and one oxygen molecule) come together to form two particles of water… or 200 particles react with 100 particles to form 200 particles… or 2 million particles react with 1 million particles to form 2 million particles… etc. The observable effect is that after the reaction, when all of the hydrogen and oxygen gases have become H20 gas, the volume of gas falls to two-thirds of the starting volume.

As a result of these observations Avogadro became the first scientist to realize that elements could exist as molecules rather than as individual atoms. For example, he recognized that the oxygen around us exists as a molecule in which two atoms of oxygen are linked. Other scientists in the field, such as Dalton, believed that only compounds could form molecules while all elements existed as single atoms. Avogadro realized that elements could exist in the form of molecules with individual atoms joined together. In 1815 Avogadro published a further paper in Journal de Physique discussing the masses of atoms, their compounds and their gas densities. In 1821, now writing as the professor of mathematical physics at the University of Turin, he published a further paper looking at the masses of atoms and the proportions in which they combine. Between 1837 and 1841 Avogadro published four weighty volumes looking in detail at the physics of matter. Avogadro’s findings were almost completely ignored until Stanislao Cannizarro presented them at the Karlsruhe Conference in 1860, four years after Avogadro’s death. This conference had been called to remedy the scientific confusion that existed about atoms, molecules and their masses.

Even after Cannizarro presented his work not all scientists agreed with it. Another decade passed – with continued strong advocacy from Cannizarro – before Avogardo’s hypothesis became more widely accepted and became Avogadro’s Law. Today Avogadro is regarded as one of the founders of atomic-molecular chemistry. Why Was Avogadro’s Hypothesis Rejected For Decades? There are a number of reasons why Avogadro’s work was not accepted quickly: • He published his work in Journal de Physique, which was not a very well read journal. • The theories of better known scientists of the time, like John Dalton and Jöns Jacob Berzelius, disagreed with Avogadro’s work. • Italy – the country of Leonardo da Vinci and Galileo – was no longer regarded as a country where great science was done. Realizing this, a few years earlier, Avogadro’s compatriot Alessandro Volta had traveled out of Italy to make himself and his work known to scientists in other countries. Avogadro stayed in Italy and did not make personal contact with foreign scientists to help bridge the gap as Volta had done.

Even though we like to think of science as a noble endeavor, where the truth will quickly become obvious, this is not always the case. Scientists are just people; they can be stubborn, just as we all can be at times. If you have a controversial new scientific theory, and you are not well-known, it seems that personal contact can sometimes achieve as much as a scientific paper – or even more. Avogadro’s Constant Avogadro’s constant is one of the most important numbers in chemistry. Its value is 6.02214129 x 1023. Avogadro did not calculate this number, but its existence follows logically from his hypothesis and work. Avogadro’s constant is the number of particles (atoms or molecules) in one mole of any substance. For example, 12 grams of carbon contain 6.02214129 x 1023 carbon atoms. Avogadro’s constant is an enormous number.

If you could save a million dollars a second, it would take you longer than the universe is believed to have existed to save 6.02214129 x 1023 dollars; which all goes to show that it takes a lot of atoms to make a small amount of matter! “My studies of the natural sciences have particularly involved that part of physics which looks at the atomic world: the properties of molecules, the forces involved in their movement, the heat capacity of different substances, expansion of gases by heat, and the density and pressure of gases.” Amedeo Avogadro, 1776 – 1856 The End Amedeo Avogadro died aged 79 on July 9th, 1856 in Turin. He was buried in the cemetery of Quaregna. Advertisements Author of this page: The Doc Images of Avogadro colorized and enhanced by this website. © All rights reserved. Please use the following MLA compliant citation: “Amedeo Avogadro.” Famous Scientists. famousscientists.org. 12 Jan. 2015. Web. < 7 Scientists whose ideas were rejected during… J. J.

Thomson Avogadro's Hypothesis Avogadro's Hypothesis Science creates knowledge using two processes: by making observations (to collect empirical evidence and create scientific laws ) by explaining the reason for those observations (to create and test hypotheses) Occasionally both processes are carried out by the same person; often however, the observations have occurred long before a satisfactory explanation for them is thought up. An hypothesis is a tentative (as opposed to a theory which is well tested) explanation for observed events. An hypothesis is not a prediction, but it must allow you to make predictions which can be tested by experiment. When the results of those experiments are as predicted, it lends support to the hypothesis as a good explanation, and its eventual acceptance as a theory. If the results are not as predicted, the hypothesis must be modified, or replaced with a better explanation. No statement is an hypothesis unless it suggests a cause for an effect, and unless it has the possibility of being wrong. Which of the following is not a characteristic of an hypothesis?

it lets us make predictions that can be experimentally tested like a theory, it may be revised when it is well established it will become a scientific law it suggests a cause - effect relationship to explain why some event occurs Truly revolutionary hypotheses, the kind that invalidate established beliefs, have often been initially rejected. This is especially true when a new idea contradicts that of well known scientists. Even scientists whose ideas were initially rejected accept this as the way science should work. For science is also full of examples of revolutionary ideas that did not stand up under the scrutiny of testing and verification. The non-discovery of cold fusion is probably one of the best known examples of this process in recent scientific history. Skepticism is one of the scientist's main tools. The late 1700's and early 1800's was a period of scientific revolution. Old ideas were crumbling, and new ones coming into place. It was an exciting time for chemistry. New concepts flowed out to explain the behavior of a new universe unfolding from Lavosier's combustion theory. New elements were being discovered, and new observations made that required explanation.

Among those observations was the work of Joseph Louis Gay-Lussac (1778-1850), whose own father was arrested in 1793, just one year before Lavoisier lost his head to the guillotine in the French Revolution. One of the easiest ways to measure a gas is by volume (provided that we keep the temperature and pressure the same). In 1800 two English chemists, William Nicholson and Anthony Carlise had demonstrated that water could be decomposed by an electric current into two volumes of Hydrogen gas, and one volume of Oxygen -- a simple whole number ratio of 2:1. Following up on this idea, Gay-Lussac observed that when gases combined in chemical reactions they did so in ratios, by volume, that were always simple whole numbers. The following table shows this for a few reactions. Combining Volumes of Gases Cl50O25 none of them However, this observation was not accepted by everyone. John Dalton for one, considered Gay-Lussac's work wrong. He believed water contained just one atom of hydrogen and one atom of oxygen ( in the symbols of the day). In today's symbols the forumla would be HO .

If this was the right formula for water, then both Gay-Lussac's law of combining volumes and Lavoisier's law of conservation of mass could not be right. This is illustrated in this diagram. Obviously something must be wrong. If two volumes of hydrogen react with one volume of oxygen as shown to produce Dalton's HO molecules, then there would have to be some hydrogen left over. This would be a violation of the Law of Conservation of mass, and totally contradictory to Dalton's atomic theory. Dalton's response was to declare that only one volume of hydrogen reacted, and that Gay-Lussac's work was inaccurate. Since we don't think water is HO today, Dalton's idea must have been all wet! His conclusion that Gay-Lussac was wrong was not agreed to by everyone , including an Italian chemist, Amedeo Avogadro. Lorenzo Romano Amedeo Carlo Avogadro di Quareqa e di Carreto - Avogadro for short - was born in Turin, Italy in 1776, into a family of church lawyers. He too studied ecclesiastical law, earning his Bachelor's degree at the age of 16, and a Doctorate by 20.

After three years he decided that a career in law was not to his bent, and he began to study science and mathematics in earnest. In 1809 he was appointed professor of Natural Philosopy at the Royal College of Vercelli in isolated Northern Italy. In 1811 he proposed his now famous hypothesis that equal volumes of gases, at the same temperature and pressure, contain equal numbers of molecules. The reaction between Hydrogen and Oxygen according to Avogadro is shown in the following diagram. Notice that each container has the same number of molecules in it. The ratio of the volumes is 2 hydrogen to 1 oxygen to 2 water, and there is nothing left over. This proposal, published in the French Journal de Physique went neglected by the European scientific community. Partly because it was published in an obscure journal, partly because it totally contradicted the opinions of John Dalton, it took fifty years before Avogadro's work was recognized. Avogadro made the distinction between atoms and molecules, which today seems clear. However, Dalton rejected Avogadro's hypothesis because Dalton believed that atoms of the same kind could not combine.

Since it was believed that atoms were held together by an electrical force, only unlike atoms would be attracted together, and like atoms should repel. Therefore it seemed impossible for a molecule of oxygen, O2, to exist. Avogadro's work, even if it was read appears not to have been understood, and was pushed into the dark recesses of chemistry libraries and ignored. Avogadro continued to teach at the university of Turin, when it was not closed because of the political upheavals going on in Italy at the time, and died in 1854, an unknown figure. Which of the following statements is true? Equal volumes of gases contain equal numbers of particles Molecules such as O2 cannot exist because their like particles would repel According to Dalton the formula for ammonia would be NH When Avogadro made up his hypothesis, he was a lawyer for the Catholic church Without a clear distinction between atoms and molecules, European chemistry blundered ahead. By 1860 enough data had been collected on atomic weights to clearly indicate that something was seriously wrong. The world's first international conference on Chemistry was called at Karlsruhe, Germany in 1860 to try and resolve the issues.

At this meeting, Stanislao Cannizaro , Avogadro's countryman, younger by fifty years, made a strong speech in which he explained the ideas of Avogadro that he had discovered two years earlier. He distributed copies of his " Sketch of a Course of Chemical Philosophy " which clearly showed the importance of Avogadro's principle, resolving the confusion with atomic and molecular weights. Very shortly, Avogadro's work was recognized for its brilliant deduction. With a clear understanding of how to measure atomic and molecular weights established, Chemistry proceeded rapidly. Less than a decade later, Dmitri Mendeleev who had also attended the Karlsruhe conference was able to use the newly established atomic weights to develop his periodic table. While chemistry is generally agreed to have started with Lavoisier, it could well be argued that modern chemistry started after the Karlruhe conference and the acceptance of Avogadro's hypothesis. In recognition of his idea, the number of fundamental particles in a mole of substance is called Avogadro's number.

Although Avogadro had no idea what the number of particles in equal volumes of gases was, and did nothing to measure it, his hypothesis did lead to the eventual determination of this number, 6.02 x 1023 How was the Avogadro's law discovered? | Socratic How was the Avogadro's law discovered? Start with a one sentence answer Then teach the underlying concepts Don't copy without citing sources Write a one sentence answer... Answer: I want someone to double check my answer Describe your changes (optional) 200 Cancel Sep 19, 2014 It comes from a brilliant insight in 1811 by the Italian chemist Amedeo Avogadro. In 1805, John Dalton proposed that atoms of different elements combine in simple whole-number ratios to form chemical compounds. In 1808, Joseph Louis Gay-Lussac reported that, when gases react to form other gases, their volumes are in whole number ratios. At the time, there was much confusion about atoms and molecules. Avogadro reasoned that simple gases were not formed of solitary atoms. Instead, they consisted of molecules containing two or more atoms.

This resolved much of the confusion that then existed. In 1811, Avogadro brilliantly hypothesized that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. Everything fell into place. Avogadro could now explain Gay-Lussac's observation that the volume of water vapour was twice the volume of the oxygen used to form it. According to Avogadro, the molecule of oxygen had split into two atoms as it formed the water vapour. Avogadro’s findings were almost completely neglected for almost fifty years. In 1860, the Italian chemist Stanislao Cannizzaro forcefully presented Avogadro's ideas during the atomic weight deliberations of the Karlsruhe Congress. He demonstrated convincingly that chemists could use Avogadro’s Law to determine molar masses and atomic masses. The scientific community finally accepted Avogadro's Law. Was this helpful? Let the contributor know! Yes

Question: Who, in 1811, postulated that 'when at the same temperature and pressure, equal volumes of gases contain equal numbers of molecules'?
Answer: Avogadro