Atom

VII

History of Atomic Theory

Beginning with Democritus, who lived during the late 5th and early 4th centuries bc, Greek philosophers developed a theory of matter that was not based on experimental evidence, but on their attempts to understand the universe in philosophical terms. According to this theory, all matter was composed of tiny, indivisible particles called atoms (from the Greek word atomos, meaning “indivisible”). If a sample of a pure element was divided into smaller and smaller parts, eventually a point would be reached at which no further cutting would be possible—this was the atom of that element, the smallest possible bit of that element.

According to the ancient Greeks, atoms were all made of the same basic material, but atoms of different elements had different sizes and shapes. The sizes, shapes, and arrangements of a material’s atoms determined the material’s properties. For example, the atoms of a fluid were smooth so that they could easily slide over one another, while the atoms of a solid were rough and jagged so that they could attach to one another. Other than the atoms, matter was empty space. Atoms and empty space were believed to be the ultimate reality.

Although the notion of atoms as tiny bits of elemental matter is consistent with modern atomic theory, the researchers of prior eras did not understand the nature of atoms or their interactions in materials. For centuries scientists did not have the methods or technology to test their theories about the basic structure of matter, so people accepted the ancient Greek view.

A

The Birth of the Modern Atomic Theory

The work of British chemist John Dalton at the beginning of the 19th century revealed some of the first clues about the true nature of atoms. Dalton studied how quantities of different elements, such as hydrogen and oxygen, could combine to make other substances, such as water. In his book A New System of Chemical Philosophy (1808), Dalton made two assertions about atoms: (1) atoms of each element are all identical to one another but different from the atoms of all other elements, and (2) atoms of different elements can combine to form more complex substances.

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Dalton’s idea that different elements had different atoms was unlike the Greek idea of atoms. The characteristics of Dalton’s atoms determined the chemical and physical properties of a substance, no matter what the substance’s form. For example, carbon atoms can form both hard diamonds and soft graphite. In the Greek theory of atoms, diamond atoms would be very different from graphite atoms. In Dalton’s theory, diamond atoms would be very similar to graphite atoms because both substances are composed of the same chemical element.

While developing his theory of atoms, Dalton observed that two elements can combine in more than one way. For example, modern scientists know that carbon monoxide (CO) and carbon dioxide (CO₂) are both compounds of carbon and oxygen. According to Dalton’s experiments, the quantities of an element needed to form different compounds are always whole-number multiples of one another. For example, two times as much oxygen is needed to form a liter of CO₂ than is needed to form a liter of CO. Dalton correctly concluded that compounds were created when atoms of pure elements joined together in fixed proportions to form units that scientists today call molecules.

A

1

States of Matter

Scientists in the early 19th century struggled in another area of atomic theory. They tried to understand how atoms of a single element could exist in solid, liquid, and gaseous forms. Scientists correctly proposed that atoms in a solid attract each other with enough force to hold the solid together, but they did not understand why the atoms of liquids and gases did not attract each other as strongly. Some scientists theorized that the forces between atoms were attractive at short distances (such as when the atoms were packed very close together to form a solid) and repulsive at larger distances (such as in a gas, where the atoms are on the average relatively far apart).

Scientists had difficulty solving the problem of states of matter because they did not adequately understand the nature of heat. Today scientists recognize that heat is a form of energy, and that different amounts of this energy in a substance lead to different states of matter. In the 19th century, however, people believed that heat was a material substance, called caloric, that could be transferred from one object to another. This explanation of heat was called the caloric theory. Dalton used the caloric theory to propose that each molecule of a gas is surrounded by caloric, which exerts a repulsive force on other molecules. According to Dalton’s theory, as a gas is heated, more caloric is added to the gas, which increases the repulsive force between the molecules. More caloric would also cause the gas to exert a greater pressure on the walls of its container, in accordance with scientists’ experiments.

This early explanation of heat and states of matter broke down when experiments in the middle of the 19th century showed that heat could change into energy of motion. The laws of physics state that the amount of energy in a system cannot increase, so scientists had to accept that heat must be energy, not a substance. This revelation required a new theory of how atoms in different states of matter behave.

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2

Behavior of Gases

In the early 19th century Italian chemist Amedeo Avogadro made an important advance in the understanding of how atoms and molecules in a gas behave. Avogadro began his work from a theory developed by Dalton. Dalton’s theory proposed that a gaseous compound, formed by combining equal numbers of atoms of two elements, should have the same number of molecules as the atoms in one of the original elements. For example, ten atoms of the element hydrogen (H) combine with ten atoms of chlorine (Cl) to form ten gaseous hydrogen chloride (HCl) molecules.

In 1811 Avogadro developed a law of physics that seemed to contradict Dalton’s theory. Avogadro’s law states that equal volumes of different gases contain the same number of particles (atoms or molecules) if both gases are at the same temperature and pressure. In Dalton’s experiment, the volume of the original vessels containing the hydrogen or chlorine gases was the same as the volume of the vessel containing the hydrogen chloride gas. The pressures of the original hydrogen and chlorine gases were equal, but the pressure of the hydrochloric gas was twice as great as either of the original gases. According to Avogadro’s law, this doubled pressure would mean that there were twice as many hydrogen chloride gas particles than there had been chlorine particles prior to their combination.

To reconcile the results of Dalton’s experiment with his new rule, Avogadro was forced to conclude that the original vessels of hydrogen or chlorine contained only half as many particles as Dalton had thought. Dalton, however, knew the total weight of each gas in the vessels, as well as the weight of an individual atom of each gas, so he knew the total number of atoms of each gas that was present in the vessels. Avogadro reconciled the fact that there were twice as many atoms as there were particles in the vessels by proposing that gases such as hydrogen and chlorine are really made up of molecules of hydrogen and chlorine, with two atoms in each molecule. Today scientists write the chemical symbols for hydrogen and chlorine as H₂ and Cl_2, respectively, indicating that there are two atoms in each molecule. One molecule of hydrogen and one molecule of chlorine combine to form two molecules of hydrogen chlorine (H₂ + Cl₂ → 2HCl). The sample of hydrogen chloride contains twice the number of particles as either the hydrogen or chlorine because two molecules of hydrogen chloride form when a molecule of hydrogen combines with a molecule of chlorine.

B

Electrical Forces in Atoms

The work of Dalton and Avogadro led to a consistent view of the quantities of different gases that could be combined to form compounds, but scientists still did not understand the nature of the forces that attracted the atoms to one another in compounds and molecules. Scientists suspected that electrical forces might have something to do with that attraction, but they found it difficult to understand how electrical forces could allow two identical, neutral hydrogen atoms to attract one another to form a hydrogen molecule.

In the 1830s, British physicist Michael Faraday took the first significant step toward appreciating the importance of electrical forces in compounds. Faraday placed two electrodes connected to opposite terminals of a battery into a solution of water containing a dissolved compound. As the electric current flowed through the solution, Faraday observed that one of the elements that comprised the dissolved compound became deposited on one electrode while the other element became deposited on the other electrode. The electric current provided by the electrodes undid the coupling of atoms in the compound. Faraday also observed that the quantity of each element deposited on an electrode was directly proportional to the total quantity of current that flowed through the solution—the stronger the current, the more material became deposited on the electrode. This discovery made it clear that electrical forces must be in some way responsible for the joining of atoms in compounds.

Despite these significant discoveries, most scientists did not immediately accept that atoms as described by Dalton, Faraday, and Avogadro were responsible for the chemical and physical behavior of substances. Before the end of the 19th century, many scientists believed that all chemical and physical properties could be determined by the rules of heat, an understanding of atoms closer to that of the Greek philosophers. The development of the science of thermodynamics (the scientific study of heat) and the recognition that heat was a form of energy eliminated the role of caloric in atomic theory and made atomic theory more acceptable. The new theory of heat, called the kinetic theory, said that the atoms or molecules of a substance move faster, or gain kinetic energy, as heat energy is added to the substance. Nevertheless, a small but powerful group of scientists still did not accept the existence of atoms—they regarded atoms as convenient mathematical devices that explained the chemistry of compounds, not as real entities.

In 1905 French chemist Jean-Baptiste Perrin performed the final experiments that helped prove the atomic theory of matter. Perrin observed the irregular wiggling of pollen grains suspended in a liquid (a phenomenon called Brownian motion) and correctly explained that the wiggling was the result of atoms of the fluid colliding with the pollen grains. This experiment showed that the idea that materials were composed of real atoms in thermal motion was in fact correct.

As scientists began to accept atomic theory, researchers turned their efforts to understanding the electrical properties of the atom. Several scientists, most notably British scientist Sir William Crookes, studied the effects of sending electric current through a gas. The scientists placed a very small amount of gas in a sealed glass tube. The tube had electrodes at either end. When an electric current was applied to the gas, a stream of electrically charged particles flowed from one of the electrodes. This electrode was called the cathode, and the particles were called cathode rays.

At first scientists believed that the rays were composed of charged atoms or molecules, but experiments showed that the cathode rays could penetrate thin sheets of material, which would not be possible for a particle as large as an atom or a molecule. British physicist Sir Joseph John Thomson measured the velocity of the cathode rays and showed that they were much too fast to be atoms or molecules. No known force could accelerate a particle as heavy as an atom or a molecule to such a high speed. Thomson also measured the ratio of the charge of a cathode ray to the mass of the cathode ray. The value he measured was about 1,000 times larger than any previous measurement associated with charged atoms or molecules, indicating that within cathode rays particularly tiny masses carried relatively large amounts of charge. Thomson studied different gases and always found the same value for the charge-to-mass ratio. He concluded that he was observing a new type of particle, which carried a negative electric charge but was about a thousand times less massive than the lightest known atom. He also concluded that these particles were constituents of all atoms. Today scientists know these particles as electrons, and Thomson is credited with their discovery.