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An atom is the smallest possible unit of an element. Since all forms of matter consist of a combination of one or more elements, atoms are the building blocks that constitute all the matter in the universe. Currently, 110 different elements, and thus 110 different kinds of atoms, are known to exist.

Our current understanding of the nature of atoms has evolved from the ancient, untested ideas of Greek philosophers, partly as a result of modern technology that has produced images of atoms.

The Greek Atomistic Philosophy

The earliest ideas concerning atoms can be traced to the Greek philosophers, who pursued wisdom, knowledge, and truth through argument and reason. Greek scientific theories were largely based on speculation, sometimes based on observations of natural phenomena and sometimes not. The idea of designing and performing experiments rarely occurred to Greek philosophers, to whom abstract intellectual activity was the only worthy pastime.

Empedocles, a Greek philosopher active around 450 B.C. , proposed that there were four fundamental substances—earth, air, fire, and water—which, in various proportions, constituted all matter. Empedocles, thus, formulated the idea of an elemental substance, a substance that is the ultimate constituent of matter; the chemical elements are modern science's fundamental substances. An atomic theory of matter was proposed by Leucippus, another Greek philosopher, around 478 B.C. Our knowledge of the atomic theory of Leucippus is derived almost entirely from the writings of his student, Democritus, who lived around 420 B.C. Democritus maintained that all materials in the world were composed of atoms (from the Greek atomos , meaning indivisible). According to Democritus, atoms of different shapes, arranged and positioned differently relative to each other, accounted for the different materials of the world. Atoms were supposed to be in random perpetual motion in a void; that is, in nothingness. According to Democritus, the feel and taste of a substance was thought to be the effect of the atoms of the substance on the atoms of our sense organs. The atomic theory of Democritus provided the basis for an explanation of the changes that occur when matter is chemically transformed. Unfortunately, the theory was rejected by Aristotle (384–322 B.C. ) who became the most powerful and famous of the Greek scientific philosophers. However, Aristotle adopted and developed Empedocles's ideas of elemental substances. Aristotle's elemental ideas are summarized in a diagram (shown in Figure 1), which associated the four elemental substances with four qualities: hot, moist, cold, and dry. Earth was dry and cold; water was cold and moist; air was moist and hot; and fire was hot and dry. Every substance was composed of combinations of the four elements, and changes (which we now call chemical ) were explained by an alteration in the proportions of the four elements. One element could be converted into the other by the addition or removal of the appropriate qualities. There were, essentially, no attempts to produce evidence to support this four-element theory, and, since Aristotle's scientific philosophy held sway for 2,000 years, there was no progress in the development of the atomic concept. The tenuous relationship between elements and atoms had been severed when Aristotle rejected the ideas of Democritus. Had the Greek philosophers been open to the idea of experimentation, atomic theory, indeed all of science, could have progressed more rapidly.

Figure 1. Aristotle's four-element diagram.
Figure 1. Aristotle's four-element diagram.

The Rise of Experimentation

The basis of modern science began to emerge in the seventeenth century, which is often recognized as the beginning of the Scientific Revolution. Conceptually, the Scientific Revolution can be thought of as a battle between three different ways of looking at the natural world: the Aristotelian, the magical, and the mechanical. The seventeenth century saw the rise of experimental science. The idea of making observations was not new. However, Sir Francis Bacon (1561–1626) emphasized that experiments should be planned and the results carefully recorded so they could be repeated and verified, an attitude that infuses the core idea of modern science. Among the early experimentalists was Robert Boyle (1627–1691), who studied quantitatively the compression and expansion of air, which led him to the idea that air was composed of particles that he called corpuscles , which he maintained were in constant motion. Boyle's description of corpuscular motion presages the kinetic molecular theory.

The Chemical Atom

An atomic theory based on chemical concepts began to emerge from the work of Antoine Lavoisier (1743–1794), whose careful quantitative experiments led to an operational definition of an element: An element was a substance that could not be decomposed by chemical processes. In other words, if a chemist could not decompose a substance, it must be an element. This point of view obviously put a premium on the ability of chemists to manipulate substances. Inspection of Lavoisier's list of elements, published in 1789, shows a number of substances, such as silica (SiO 2 ), alumina (Al 2 O 3 ), and baryta (BaO), which today are recognized as very stable compounds. The chemists of Lavoisier's time simply did not have the tools to decompose these substances further to silicon, aluminum, and barium, respectively. The composition of all compounds could be expressed in terms of the elemental substances, but it was the quantitative mass relationship of compounds that was the key to deducing the reality of the chemical atom.

Lavoisier's successful use of precise mass measurements essentially launched the field of analytical chemistry, which was thoroughly developed by Martin Klaproth (1743–1817). Lavoisier established the concept of mass conservation in chemical reactions, and, late in the eighteenth century, there was a general acceptance of the concept of definite proportions (constant composition) in chemical compounds, but not without controversy. Claude-Louis Berthollet (1748–1822) maintained that the composition of compounds could be variable, citing, for example, analytical results on the oxides of copper, which gave a variety of results, depending on the method of synthesis . Joseph-Louis Proust (1754–1826), over a period of eight years, showed that the variable compositions, even with very accurate analytical data, were due to the formation of different mixtures of two oxides of copper, CuO and Cu 2 O. Each oxide obeyed the law of constant composition, but reactions that were supposed to lead to "copper oxide" often produced mixtures, the proportions of which depended on the conditions of the reaction. Proust's proof of the law of constant composition was important, because compounds with variable composition could not be accommodated within the evolving chemical atomic theory.


Little is known for certain about Democritus of Abbera (c.460 B.C.E. –c.362 B.C.E. ). None of his writings has survived intact. It is known from others (both students and detractors) that Democritus was one of the earliest advocates of a theory that all matter exists as collections of very small, permanent, indivisible particle called atoms.

—David A. Bassett

John Dalton (1766–1844), a self-educated English scientist, was primarily interested in meteorology and is credited with being the first to describe color blindness, a condition with which he was burdened throughout his life. Color blindness is a disadvantage for a chemist, who must be able to see color changes when working with chemicals. Some have suggested that his affliction was one reason why Dalton was a rather clumsy and slip-shod experimenter. Gaseous behavior had been well established, starting with the experiments of Boyle. Dalton could not help supposing, as others previously did, that gaseous matter was composed of particles. But Dalton took the next and, ultimately, most important steps in assuming that all matter—gaseous, liquid, and solid—consists of these small particles. The law of definite proportions (constant composition) as articulated by Proust, suggested to Dalton that a compound might contain two elements in the ratio of, for example, 4 to 1, but never 4.1 to 1 or 3.9 to 1. This observation could easily be explained by supposing that each element was made up of individual particles.

Dalton's atomic theory can be succinctly summarized by the following statements:

Elements are composed of extremely small particles called atoms.

All atoms of a given element have identical properties, and those properties differ from those of other elements.

Compounds are formed when atoms of different elements combine with one another in small whole numbers.

The relative numbers and kinds of atoms are constant in a given compound.

Dalton recognized the similarity of his theory to that of Democritus, advanced twenty-one centuries earlier when the Greek philosopher called these small particles atoms , and, presumably, implied by using that word that these particles were indivisible. In Dalton's representation (Figure 2) the elements were shown as small spheres, each with a separate identity. Compounds of elements were shown by combining the appropriate elemental representations in the correct proportions, to produce complex symbols that seem to echo our present use of standard chemical formulas. Dalton's symbols—circles with increasingly complex inserts and decorations—were not adopted by the chemical community. Current chemical symbols (formulas) are derived from the suggestions of Jöns Berzelius (1779–1848). Berzelius also chose oxygen to be the standard reference for atomic mass (O = 16.00 AMU). Berzelius produced a list of atomic masses that were much closer to those that are currently accepted because he had developed a better way to obtain the formulas of substances. Whereas Dalton assumed that water had the formula HO, Berzelius showed it to be H 2 O. The property of atoms of interest to Dalton were their relative masses, and Dalton produced a table of atomic masses (Table 1) that was seriously deficient because he did not appreciate that atoms did not have to be in a one-to-one ratio; using more modern ideas, Dalton assumed, incorrectly, that all atoms had a valence of one (1). Thus, if the atomic mass of hydrogen is arbitrarily assigned to be 1, the atomic mass of oxygen is 8 on the Dalton scale. Dalton, of course, was wrong, because a water molecule contains two atoms of hydrogen for every oxygen atom, so that the individual oxygen atom is eight times as heavy as two hydrogen atoms or sixteen times as heavy as a single hydrogen atom. There was no way that Dalton could have known, from the data available, that the formula for water is H 2 O.

Figure 2. Dalton's atomic symbols are described as "simple." The increasingly complex combination of symbols represent binary, ternary, quaternary, etc., compounds. Thus, Number 28 is a compound atom of carbonic acid (carbon dioxide), and number 31 is a compound atom of sulphuric acid (sulphyr trioxide).
Figure 2. Dalton's atomic symbols are described as "simple." The increasingly complex combination of symbols represent binary, ternary, quaternary, etc., compounds. Thus, Number 28 is a compound atom of carbonic acid (carbon dioxide), and number 31 is a compound atom of sulphuric acid (sulphyr trioxide).

Dalton's atomic theory explained the law of multiple proportions. For example, it is known that mercury forms two oxides: a black substance containing 3.8 percent oxygen and 96.2 percent mercury, and a red compound containing 7.4 percent oxygen and 92.6 percent mercury. Dalton's theory states that the atoms of mercury (Hg) and oxygen (O) must combine in whole numbers, so the two compounds might be HgO and Hg 2 O, for example. Furthermore, Dalton's theory states that each element has a characteristic mass—perhaps 9 mass units for Hg and 4 mass units for O (the

Table 1. Dalton's first set of atomic weight values (1805).
Table 1. Dalton's first set of atomic weight values (1805).

Hydrogen 1
Azot 4.2
Carbon 4.3
Ammonia 5.2
Oxygen 5.5
Water 6.5
Phosphorus 7.2
Phosphuretted hydrogen 8.2
Nitrous gas 9.3
Ether 9.6
Gaseous oxide of carbon 9.8
Nitrous oxide 13.7
Sulphur 14.4
Nitric acid 15.2
Sulphuretted hydrogen 15.4
Carbonic acid 15.3
Alcohol 15.1
Sulphureous acid 19.9
Sulphuric acid 25.4
Carburetted hydrogen from stagnant water 6.3
Olefiant gas 5.3

numbers were chosen arbitrarily, here). Given these assumptions, the relevant concepts are shown in Table 2.

The assumed formulas are presented in line 1. The percent composition of each compound, calculated in the usual way, is presented in line 3, showing that these two compounds, indeed, have different compositions, as required by the law of multiple proportions. Line 4 contains the ratio of the mass of mercury to the mass of oxygen, for each compound. Those ratios can be expressed as the ratio of simple whole numbers (2.25:4.5 = 1:2), fulfilling a condition required by the law of multiple proportions. Notice that Dalton's ideas do not depend upon the values assigned to the elements or the formulas for the compounds involved. Indeed, the question as to which compound, red or black, is associated with which formula cannot be answered from the data available. Thus, although Dalton was unable to establish an atomic mass scale, his general theory did provide an understanding of the three mass-related laws: conservation, constant composition, and multiple proportion. Other information was required to establish the relative masses of atoms.

The other piece of the puzzle of relative atomic masses was provided by Joseph-Louis Gay-Lussac (1778–1850), who published a paper on volume relationships in reactions of gases. Gay-Lussac made no attempt to interpret his results, and Dalton questioned the paper's validity, not realizing that the law of combining volumes was really a verification of his atomic theory! Gay-Lussac's experiments revealed, for example, that 2 volumes of carbon monoxide combine with 1 volume of oxygen to form 2 volumes of carbon dioxide. Reactions of other gaseous substances showed similar volume relationships. Gay-Lussac's law of combining volumes suggested, clearly, that equal volumes of different gases under similar conditions of temperature and pressure contain the same number of reactive particles (molecules). Thus, if 1 volume of ammonia gas (NH 3 ) combines exactly with 1 volume of hydrogen chloride gas (HCl) to form a salt (NH 4 Cl), it is natural to conclude that each volume of gas must contain the same number of particles.

Table 2. Law of multiple proportions.
Table 2. Law of multiple proportions.

Assumed formula HgO Hg 2 O
Total mass of compound 9 + 4 = 13 9 + 9 + 4 = 22
% composition % Hg 69.2; % O = 30.8 % Hg = 81.8; % O = 18.2
Mass Hg/Mass O 9/4 = 2.25 18/4 = 4.5

At least one of the implications of Gay-Lussac's law was troubling to the chemistry community. For example, in the formation of water, 2 volumes of hydrogen gas combined with 1 volume of oxygen gas to produce 2 volumes of steam (water in the gaseous state). These observations produced, at the time, an apparent puzzle. If each volume of gas contains n particles (molecules), 2 volumes of steam must contain 2 n particles. Now, if each water particle contains at least 1 oxygen atom, how is it possible to get two oxygen atoms (corresponding to 2 n water molecules) from n oxygen particles? The obvious answer to this question is that each oxygen particle contains two oxygen atoms. This is equivalent to stating that the oxygen molecule consists of two oxygen atoms , or that oxygen gas is diatomic (O 2 ). Amedeo Avogadro (1776–1856) an Italian physicist, resolved the problem by adopting the hypothesis that equal volumes of gases under the same conditions contain equal numbers of particles (molecules). His terminology for what we now call an atom of, for instance, oxygen, was half molecule. Similar reasoning involving the combining of volumes of hydrogen and oxygen to form steam leads to the conclusion that hydrogen gas is also diatomic (H 2 ). Despite the soundness of Avogadro's reasoning, his hypothesis was generally rejected or ignored. Dalton never appreciated its significance because he refused to accept the experimental validity of Gay-Lussac's law.

Avogadro's hypothesis—equal volumes of gases contain equal numbers of particles—lay dormant for nearly a half-century, until 1860 when a general meeting of chemists assembled in Karlsruhe, Germany, to address conceptual problems associated with determining the atomic masses of the elements. Two years earlier, Stanislao Cannizzaro (1826–1910) had published a paper in which, using Avogadro's hypothesis and vapor density data, he was able to establish a scale of relative atomic masses of the elements. The paper, when it was published, was generally ignored, but its contents became the focal point of the Karlsruhe Conference.

Cannizzaro's argument can be easily demonstrated using the compounds hydrogen chloride, water, ammonia, and methane, and the element hydrogen, which had been shown to be diatomic (H 2 ) by using Gay-Lussac's reasoning and his law of combining volumes. The experimental values for vapor density of these substances, all determined under the same conditions of temperature and pressure, are also required for Cannizzaro's method for establishing atomic masses. The relevant information is gathered in Table 3. The densities of these gaseous substances (at 100° C and one atmosphere pressure) are expressed in grams per liter. The masses of the substances (in one liter) are the masses of equal numbers of molecules of each substance; the specific number of molecules is unknown, of course, but that number is unnecessary for the Cannizzaro analysis. If that unknown number of molecules is called N , and if m H represents the mass of a single hydrogen atom, then m H × 2 N is the total

Table 3. Cannizzaro's method of molecular mass determination.
Table 3. Cannizzaro's method of molecular mass determination.

Gaseous Substance Density g/L 1 Relative to Mass of an H Atom (Molecular Mass, Relative to H = 1) % Hydrogen Relative Mass of H Present Number H Atoms Present in a Molecule Formula Mass of "Other" Atoms
1 Density reported for conditions of 100°C and one atmosphere pressure
Hydrogen 0.0659 2.00 100 2.00 2 H 2 H = 1
Hydrogen chloride 1.19 36.12 2.76 1.00 1 HCl Cl = 35.2
Water 0.589 17.88 11.2 2.00 2 H 2 O O = 15,88
Ammonia 0.557 16.90 17.7 3.00 3 NH 3 N = 13.90
Methane 0.524 15.90 25.1 4.00 4 CH 4 C = 11.90

mass of the hydrogen atoms in the 1 liter sample of hydrogen molecules; recall that hydrogen was shown to be diatomic (H 2 ) by Gay-Lussac's law. From this point of view, the relative masses of the molecules fall in the order of the masses in 1 liter (or their densities). The mass of the hydrogen atom was taken as the reference (H = 1) for the relative atomic masses of the elements. Thus, the mass of all the hydrogen chloride molecules in the one liter sample is m HCl N , and the ratio of the mass of a hydrogen chloride molecule to a hydrogen atom is given by:

That is, if the mass of a hydrogen atom is taken to be 1 unit of mass, the mass of the hydrogen chloride molecule is 36.12 units. All the molecular masses listed in column 3 of the table can be established in the same way—twice the ratio of the density of the molecule in question to the density of hydrogen. Using experimental analytical data (column 4), Cannizzaro was able to establish the relative mass of hydrogen in each molecule (column 5), which gave the number of hydrogen atoms present in each molecule of interest (column 6), which, in turn, produced the formula of the molecule (column 7); analytical data also quantitatively indicate the identity of the other atom in the molecule. Thus, analysis would tell us that, for example, methane contains hydrogen and carbon. Knowing the total mass of the molecule (column 3) and the mass of all the hydrogen atoms present, the mass of the "other atom" in the molecule can be established as the difference between these numbers (column 8). Thus, if the mass of the HCl molecule is 36.12 and one atom of hydrogen of mass 1.00 is present, the mass of a Cl atom is 35.12. Relative mass units are called atomic mass units , AMUs.

This very convincing use of Gay-Lussac's law and Avogadro's hypothesis by Cannizzaro quickly provided the chemical community with a direct way of establishing not only the molecular formulas of binary compounds but also the relative atomic masses of elements, starting with quantitative analytical data and the density of the appropriate gaseous substances.

The long struggle to establish the concept of the chemical atom involved many scientists working in different countries using different kinds of equipment to obtain self-consistent data. All were infused with ideas of Sir Francis Bacon, who defined the classic paradigm of experimental science—results that are derived from careful observations and that are openly reported for verification. However, not all chemists equally embraced these ideas, which were to become fundamental to their craft. For example, the great physical chemist and Nobel Prize winner Friedrich Wilhelm Ostwald (1853–1932) refused to accept the existence of atoms well into the twentieth century. Ostwald held a strong personal belief that chemists ought to confine their studies to measurable phenomena such as energy changes. The atomic theory was to Ostwald nothing more than a convenient fiction.

There are, of course, other lines of observations and arguments that lead to the conclusion that matter is particulate and, subsequently, to an ultimate atomic description of matter. One of these involves the Brownian motion of very small particles. Robert Brown (1773–1858), a Scottish botanist, observed in 1827 that individual grains of plant pollen suspended in water moved erratically. This irregular movement of individual particles of a suspension as observed with a microscope is called Brownian motion. Initially, Brown believed that this motion was caused by the "hidden life" within the pollen grains, but further studies showed that even nonliving suspensions behave in the same way. In 1905 Albert Einstein (1879–1955) worked out a mathematical analysis of Brownian motion. Einstein showed that if the water in which the particles were suspended was composed of molecules in random motion according to the requirements of the kinetic molecular theory, then the suspended particles would exhibit a random "jiggling motion" arising from the occasional uneven transfer of momentum as a result of water molecules striking the pollen grains. One might expect that the forces of the water molecules striking the pollen grains from all directions would average out to a zero net force. But Einstein showed that, occasionally, more water molecules would strike one side of a pollen grain than the other side, resulting in a movement of the pollen grain. The interesting point in Einstein's analysis is that even if each collision between a water molecule and a pollen grain transfers a minuscule amount of momentum, the enormous

Photomicrograph of atoms in a tungsten crystal, magnified 2,700,000 times.
Photomicrograph of atoms in a tungsten crystal, magnified 2,700,000 times.

number of molecules striking the pollen grain is sufficient to overcome the large momentum advantage of the pollen grain (because of its considerably larger mass than that of a water molecule). Although the Swedish chemist Theodor Svedberg (1884–1971) suggested the general molecular explanation earlier, it was Einstein who worked out the mathematical details. Einstein's analysis of Brownian motion was partially dependent on the size of the water molecules. Three years later, Jean-Baptiste Perrin (1870–1942) set about to determine the size of the water molecules from precise experimental observations of Brownian motion. In other words, Perrin assumed Einstein's equations were correct, and he made measurements of the particles' motions, which Brown had described only qualitatively. The data Perrin collected allowed him to calculate the size of water molecules. Ostwald finally yielded in his objection to the existence of atoms because Perrin had a direct measure of the effect of water molecules on macroscopic objects (pollen grains). Since water was composed of the elements hydrogen and oxygen, the reality of atoms had been experimentally proved in Ostwald's view of how chemistry should be pursued.

Ostwald's reluctance to accept the chemical atom as an entity would surely have yielded to the overwhelming evidence provided by scanning tunneling microscopy (STM). Although Ostwald did not live to see it, this technique provides such clear evidence of the reality of simple atoms that even he would have been convinced.

SEE ALSO Avogadro, Amedeo ; Berthollet, Claude-Louis ; Berzelius, Jöns JaKob ; Boyle, Robert ; Cannizzaro, Stanislao ; Dalton, John ; Einstein, Albert ; Gay-Lussac, Joseph-Louis ; Lavoisier, Antoine ; Ostwald, Friedrich Wilhelm ; Svedberg, Theodor ; Molecules .

J. J. Lagowski


Hartley, Harold (1971). Studies in the History of Chemistry. Oxford, U.K.: Clarendon Press.

Ihde, Aaron J. (1964). The Development of Modern Chemistry. New York: Harper and Row.

Lavoisier, Antoine; Fourier, Jean-Baptiste Joseph; and Faraday, Michael (1952). Great Books of the Western World , Vol. 45, tr. Robert Kerr and Alexander Freeman. Chicago: Encyclopedia Britannica.

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