Atomic theory

History

Describing characteristics of atoms

Applications of atomic theory

Atomic theory is the description of atoms, the smallest units of elements.

History

Beginning in about 600 BC, many Greek philosophers struggled to understand the nature of matter. Some said everything was made of water, which comes in three forms (solid ice, liquid water, and gaseous steam). Others believed that matter was made entirely of fire in ever-changing forms. Still others believed that whatever comprised matter, it must be something that could not be destroyed but only recombined into new forms. If they could see small enough things, they would find that the same “building blocks” they started with were still there. One of these philosophers was named Democritus. He imagined starting with a large piece of matter and gradually cutting it into smaller and smaller pieces, finally reaching a smallest possible piece. This tiniest building block that could no longer be cut he named atomos, Greek for “no-cut” (indivisible). The word atomos has been changed in modern times to “atom.” The atoms Democritus envisioned differed only in shape and size. In his theory, different objects looked different because of the way the atoms were arranged. Aristotle, one of the most influential philosophers of that time, believed in some kind of “smallest part” of matter, but did not believe these parts followed Democritus’s description. Aristotle said there were only four elements (earth, air, fire, water) and that these had some smallest unit that made up all matter. Aristotle’s teachings against the idea of Democritus’s atom were so powerful that the idea of the atom fell out of philosophical fashion for the next 2,000 years.

Although atomic theory was abandoned for this long period, scientific experimentation, especially in chemistry, flourished. From the Middle Ages (c. AD 1100) onward, many chemical reactions were studied. By the seventeenth century, some of these chemists began thinking about the reactions they were seeing in terms of smallest parts. They even began using the word atom again. One of the most famous chemists of the end of the eighteenth century was Antoine Lavoisier (1743–1794). His chemical experiments involved very careful weighing of all the chemicals. He reacted various substances until they were in their simplest states. He found two important factors: (1) the simplest substances, which he called elements, could not be broken down any further, and (2) these elements always reacted with each other in the same proportions. The same more complex substances he called compounds. For example, two volumes of hydrogen reacted exactly with one volume of oxygen to produce water. Water could be broken down to always give exactly two volumes of hydrogen and one volume of oxygen. Lavoisier had no explanation for these amazingly consistent results. However, his numerous and careful measurements provided a clue for another chemist named John Dalton (1766–1844).

Dalton realized that if elements were made up of atoms, a different atom for each different element, atomic theory could explain Lavoisier’s results. If two atoms of hydrogen always combined with one atom of oxygen, the resulting combination of atoms, called a molecule, would be water. Dalton published his explanation in 1803. This year is considered the beginning of modern atomic theory. Scientific experiments that followed Dalton were attempts to characterize how many elements there were, what the atoms of each element were like, how the atoms of each element were the same and how they differed, and, ultimately, whether there was anything smaller than an atom.

Describing characteristics of atoms

One of the first attributes of atoms to be described was relative atomic weight. Although a single atom

was too small to weigh, atoms could be compared to each other. The chemist Jons Berzelius (1799–1848) assumed that equal volumes of gases at the same temperature and pressure contained equal numbers of atoms. He used this idea to compare the weights of reacting gases. He was able to determine that, for example, oxygen atoms were 16 times heavier than hydrogen atoms. He made a list of these relative atomic weights for as many elements as he knew. He devised symbols for the elements by using the first letter or first two letters of their Latin names, a system still in use today. The symbol for hydrogen is H, for oxygen is O, for sodium (natrium, in Latin) is Na, and so on. The symbols also proved useful in describing how many atoms combine to form a molecule of a particular compound. For example, to show that water is made of two atoms of hydrogen and one atom of oxygen, the symbol for water is H₂ O. One oxygen atom can even combine with one other oxygen atom to produce a molecule of oxygen with the symbol O₂.

As more and more elements continued to be discovered, it became convenient to begin listing them in symbol form in a chart. In 1869, Dmitri Mendeleev (1834–1907) listed the elements in order of increasing atomic weight and grouped elements that seemed to have similar chemical reactions. For example, lithium (Li), sodium (Na), and potassium (K) are all metallic elements that burst into flame if they get wet. Similar elements were placed in the same column of his chart. Mendeleev began to see a pattern among the elements, where every eighth element on the atomic weight listing would belong to the same column. Because of this periodicity or repeating pattern, Mendeleev’s chart is called the periodic table of the elements. The table was so regular, in fact, that when there was a “hole” in the table, Mendeleev predicted that an element would eventually be discovered to fill it. For instance, there was a space in the table for an element with an atomic weight of about 72 (72 times heavier than hydrogen), but no known element of that weight. In 1886, 15 years after its prediction, the element germanium (Ge) was isolated and found to have an atomic weight of 72.3. Many more elements continued to be predicted and found in this way. However, as more elements were added to the periodic table, it was found that if some elements were placed in the correct column because of similar reactions, they did not follow the right order of increasing atomic weight. Some other

atomic characteristic was needed to order the elements properly. Many years passed before the correct property was found.

As chemistry experiments were searching for and characterizing more elements, other branches of science were making discoveries about electricity and light that were to contribute to the development of atomic theory. Michael Faraday (1791–1867) had done much work to characterize electricity; James Clerk Maxwell (1831–1879) characterized light. In the 1870s, William Crookes built an apparatus, now called a Crookes tube, to examine “rays” being given off by metals. He wanted to determine whether the rays were light or electricity based on Faraday’s and Maxwell’s descriptions of both. Crookes’s tube consisted of a glass bulb, from which most of the air had been removed, encasing two metal plates called electrodes. One electrode was called the anode and the other was called the cathode. The plates each had a wire leading outside the bulb to a source of electricity. When electricity was applied to the electrodes, rays appeared to come from the cathode. Crookes determined that these cathode rays were particles with a negative electrical charge that were being given off by the metal of the cathode plate. In 1897, J. J. Thomson (1856–1940) discovered that these negatively charged particles were coming out of the atoms and must have been present in the metal atoms to begin with. He called these negatively charged subatomic particles “electrons.” Since the electrons were negatively charged, the rest of the atom had to be positively charged. Thomson believed that the electrons were scattered in the atom like raisins in a positively-charged bread dough, or like plums in a pudding. Although Thomson’s “plum-pudding” model was not correct, it was the first attempt to show that atoms were more complex than just homogeneous spheres.

At the same time, scientists were examining other kinds of mysterious rays that were coming from the Crookes tube but did not originate at its cathode. In 1895, Wilhelm Roentgen (1845–1923) noticed that photographic plates held near a Crookes tube would become fogged by some invisible, unknown rays. Roentgen called these rays “x rays,” using “x” for unknown as is common in algebra. Roentgen also established the use of photographic plates as a way to take pictures of mysterious rays. He found that by blocking the x rays with his hand, for instance, bones would block the x rays but skin and tissue would not. Doctors still use Roentgen’s x rays for imaging the human body.

Photographic plates became standard equipment for scientists of Roentgen’s time. One of these scientists, Henri Becquerel (1852–1908), left some photographic plates in a drawer with uranium, a new element he was studying. When he removed the plates, he found that they had become fogged. Since there was nothing else in the drawer, he concluded that the uranium must have been giving off some type of ray. Becquerel showed that this radiation was not as penetrating as x rays since it could be blocked by paper. The element itself was actively producing radiation, a property referred to as radioactivity. Largely through the work of Pierre and Marie Curie (1859–1906; 1867–1934), more radioactive elements were found. The attempts to characterize the different types of radioactivity led to the next great chapter in the development of atomic theory.

In 1896, Ernest Rutherford (1871–1937), a student of J. J. Thomson, began studying radioactivity. By testing various elements and determining what kinds of materials could block the radiation from reaching a photographic plate, Rutherford concluded that there were two types of radioactivity coming from elements. He named them using the first two letters of the Greek alphabet, alpha and beta. Alpha radiation was made of positively charged particles about four times as heavy as a hydrogen atom. Beta radiation was made of negatively charged particles that seemed to be just like electrons. Rutherford decided to try an experiment using the alpha particles. He set up a piece of thin gold foil with photographic plates encircling it. He then allowed alpha particles to hit the gold. Most of the alpha particles went right through the gold foil, but a few of them did not. A few alpha particles were deflected from their straight course. A few even came straight backward. Rutherford wrote that it was as surprising as if one had fired a bullet at a piece of tissue paper only to have it bounce back. Rutherford concluded that since most of the alpha particles went through, the atoms of the gold must be mostly empty space, not Thomson’s space-filling plum-pudding. Since a few of the alpha particles were deflected, there must be a densely packed positive region in each atom, which he called the nucleus. With all the positive charge in the nucleus, the next question was the arrangement of the electrons in the atom.

In 1900, physicist Max Planck had been studying processes of light and heat, specifically trying to understand the light radiation given off by a “black-body,” an ideal cavity made by perfectly reflecting walls. This cavity was imagined as containing objects called oscillators, which absorbed and emitted light and heat. Given enough time, the radiation from such a black-body would produce a colored-light distribution called a spectrum, which depended only on the temperature of the black-body and not on what it was made of. Many scientists attempted to find a mathematical relationship that would predict how the oscillators of a black-body could produce a particular spectral distribution. Max Planck found the correct mathematical relationship. He assumed that the energy absorbed or emitted by the oscillators was always a multiple of some fundamental “packet of energy” he called a quantum. Objects that emit or absorb energy do it in discrete amounts, called quanta.

At the same time, there was a physicist working with Thomson and Rutherford named Niels Bohr. Bohr realized that the idea of a quantum of energy could explain how the electrons in the atom are arranged. He described the electrons as being “in orbit” around the nucleus like planets around the sun. Like oscillators in a black-body could not have just any energy, electrons in the atom could not have just any orbit. There were only certain distances that were allowed by the energy an electron has. If an electron of a particular atom absorbed the precisely right quantum of energy, it could move farther away from the nucleus. If an electron farther from the nucleus emitted the precisely right quantum of energy, it could move closer to the nucleus. The precisely right values differed for every element. These values could be determined by a process called atomic spectroscopy, an experimental technique that looked at the light spectrum produced by atoms. An atom was heated so that all of its electrons were moved far away from the nucleus. As they moved closer to the nucleus, the electrons would begin emitting their quanta of energy as light. The spectrum of light produced could be examined using a prism. The spectrum produced in this way did not show every possible color, but only those few that matched the energies corresponding to the electron orbit differences. Although later refined, Bohr’s “planetary model” of the atom explained atomic spectroscopy data well enough that scientists turned their attention back to the nucleus of the atom.

Rutherford, along with Frederick Soddy, continued work with radioactive elements. Soddy, in particular, noticed that as alpha and beta particles were emitted from atoms, the atoms changed in one of two ways: (1) the element became a totally different element with completely new chemical reactions, or (2) the element maintained the same chemical reactions and the same atomic spectrum, only changing in atomic weight.

Soddy called atoms of the second group (atoms of the same element with different atomic weights) isotopes. In any natural sample of an element, there may be several types of isotopes. As a result, the atomic weight of an element that was calculated by Berzelius was actually an average of all the isotope weights for that element. This was the reason that some elements did not fall into the correct order on Mendeleev’s periodic table—the average atomic weight depended on how much of each kind of isotope was present. Soddy suggested placing the elements in the periodic table by similarity of chemical reactions and then numbering them in order. The number assigned to each element in this way is called the atomic number. The atomic numbers were convenient ways to refer to elements.

Meanwhile, Thomson had continued his work with the Crookes tube. He found that not only were cathode rays of electrons produced, so were positive particles. After much painstaking work, he was able to separate the many different kinds of positive particles by weight. Based on these measurements, he was able to determine a fundamental particle, the smallest positive particle produced, called a proton. Since these were being produced by the atoms of the cathode, and since Rutherford showed that the nucleus of the atom was positive, Thomson realized that the nucleus of an atom must contain protons. A young scientist named Henry Moseley experimented with bombarding atoms of different elements with x rays. Just as in atomic spectroscopy, where heat gives electrons more energy, x rays give protons in the nucleus more energy. And just as electrons give out light of specific energies when they cool, the nucleus emits x rays of a specific energy when it “de-excites.” Moseley discovered that for every element the energy of the emitted x rays followed a simple mathematical relationship. The energy depended on the atomic number for that element, and the atomic number corresponded to the number of positive charges in the nucleus. So the correct ordering of the periodic table is by increasing number of protons in the atomic nucleus. The number of protons equals the number of electrons in a neutral atom. The electrons are responsible for the chemical reactions. Elements in the same column of the periodic table have similar arrangements of electrons with the highest energies, and this is why their reactions are similar.

Only one problem remained. Electrons had very little weight, 1/1,836 the weight of a proton. Yet the protons did not account for all of the atomic weight of an atom. It was not until 1932 that James Chadwick discovered the existence of a particle in the nucleus with no electrical charge but with a weight slightly greater than a proton. He named this particle the neutron. Neutrons are responsible for the existence of isotopes. Two atoms of the same element will have the same number of protons and electrons, but they might have different numbers of neutrons and therefore different atomic weights. Isotopes are named by stating the name of the element and then the number of protons plus neutrons in the nucleus. The sum of the protons and neutrons is called the mass number. For example, uranium-235 has 235 protons and neutrons. We can look on a periodic table to find uranium’s atomic number (92) which tells us the number of protons. Then by subtracting, we know that this isotope has 143 neutrons. There is another isotope of uranium,²³⁸U, with 92 protons and 146 neutrons. Some combinations of protons and neutrons are less stable than others. Picture trying to hold 10 bowling balls in your arms. There will be some arrangement where you might be able to manage it. Now try holding 11 or only nine. There might not be a stable arrangement, and you would drop the bowling balls. The same thing happens with protons and neutrons. Unstable arrangements spontaneously fall apart, emitting particles, until a stable structure is reached. This is how radioactivity like alpha particles is produced. Alpha particles are made of two protons and two neutrons tumbling out of an unstable nucleus.

Hydrogen has three kinds of isotopes: hydrogen,²H (deuterium), and ³H (tritium).

The atomic weights of the other elements were originally compared to hydrogen without specifying which isotope. It is also difficult to get single atoms of hydrogen because a lone hydrogen atom usually reacts with other atoms to form molecules such as H₂ or H₂ O. Therefore, a different element’s isotope was chosen for comparison. The atomic weights are now based on¹²C (carbon 12). This isotope of carbon has six protons and six neutrons in its nucleus. Carbon-12 was defined to be 12 atomic mass units. (Atomic mass units, abbreviated amu, are units used to compare the relative weights of atoms. One amu is less than 200 sextillionths of a gram.) Every other isotope of every other element is compared to this. Then the weights of a given element’s isotopes are averaged to give the atomic weights found on the periodic table.

Until this point in the story of the atom, all of the particles comprising the atom were thought of as hard, uniform spheres. Beginning in 1920, with the work of Louis de Broglie, this image changed. De Broglie showed that particles like electrons could sometimes have properties of waves. For instance, if water waves are produced by two sources, like dropping two pebbles into a pond, the waves can interfere with each other. This means that high spots add to make even higher spots. Low spots add to make even lower regions. When electrons were made to travel through a double slit, with some electrons going through one slit and some through the other, they effectively created two sources. The electrons showed this same kind of interference, producing a pattern on a collection plate. The ability of electrons and other particles to sometimes show properties of particles and sometimes of waves is called wave-particle duality. This complication to the nature of the electron meant that Bohr’s idea of a planetary atom was not quite right. The electrons do have different discrete energies, but they do not follow circular orbits. In 1925, Werner Heisenberg stated that the precise speed and location of a particle cannot, for fundamental physical reasons, both be known at the same time. The Heisenberg uncertainty principle inspired Erwin Schrödinger to devise an equation to calculate how an electron with a certain energy moves. Schrödinger’s equation describes regions in an atom where an electron with a certain energy is likely to be

KEY TERMS

Accelerator —A device that causes particles to move faster and faster.

Alpha particle —Two protons and two neutrons bound together and emitted from the nucleus during some kinds of radioactive decay.

Atomic mass —The mass of an atom relative to carbon-12 (¹²C, which has a mass of exactly 12 atomic mass units); also the mass, in grams, of an element that contains one mole of atoms.

Atomic mass unit (u or amu) —A unit used to express the mass of atoms equal to exactly one-twelfth of the mass of carbon-12.

Beta particle —One type of radioactive decay particle emitted from radioactive atomic nuclei. A beta particle is the same thing as an electron.

Electrode —A metal plate that carries electrical current.

Electron cloud —The image of an electron moving so fast that it seems to fill a region of space.

Interference —The combination of waves in which high spots combine to give even higher spots and low spots combine to give even lower spots.

Kinetic energy —The energy of a moving object.

Mass number —The sum of protons and neutrons.

Nucleus —The dense central part of an atom containing the protons and neutrons; plural is nuclei.

Orbital —The region of probability within an atom where an electron with a particular energy is likely to be.

Oscillators —Objects that can absorb or emit energy and convert it into kinetic energy.

Periodicity —Repeatability of a pattern.

Quantum —The amount of radiant energy in the different orbits of an electron around the nucleus of an atom.

Quarks —Believed to be the most fundamental units of protons and neutrons.

Uncertainty principle —Heisenberg’s statement that both the position and velocity of a particle cannot be known with equal precision at the same time.

Wave-particle duality —The ability of objects to show characteristics of both waves and particles.

but not exactly where it is. This region of probability is called an orbital. Electrons move about so fast within these orbitals that we can think of them as blurring into an electron cloud. Electrons move from one orbital into another by absorbing or emitting a quantum of energy, just as Bohr explained.

Applications of atomic theory

Early studies of radioactivity revealed that certain atomic nuclei were naturally radioactive. Some scientists wondered that if particles could come out of the nucleus, would it also be possible to force particles into the nucleus? In 1932, John D. Cockcroft (1897–1967) and Ernest Walton (1903–1995) succeeded in building a particle accelerator, a device that could make streams of charged particles move faster and faster. These fast particles, protons for example, were then aimed at a thin plate of a lighter element like lithium (Li). If a lithium atom nucleus “captures” a proton, the nucleus becomes unstable and breaks apart into two alpha particles. This technique of inducing radioactivity by bombardment with accelerated particles is still the most used method of studying nuclear structure and subatomic particles. Today, accelerators race the particles in straight lines or, to save land space, in ringed paths several miles in diameter.

The spontaneous rearrangement of the atomic nucleus always results in a release of energy in the form of kinetic motion in fast-moving neutrons. When a large nucleus falls apart to form smaller atoms, the process is called fission. When lighter atoms are forced together to produce a heavier atom, the process is called fusion. In either case, fast neutrons are released. These can transfer their kinetic energy to the surroundings, heating it. This heat can be used to boil water, producing steam to run a turbine that turns an electric generator. Fusion, the coming-together of nuclei to make heavier nuclei, is the process of releasing energy at the center of the sun and other stars. So much energy can be released quickly by either fission or fusion that these processes have made possible the manufacture of atomic weapons. Fusion is not yet controlled enough for running a power plant. Research continues to find a controlled method of using fusion energy, but the problem appears extremely difficult and affordable fusion power may never be achieved.

The first atomic (fission) bomb was detonated as part of the test code-named Trinity in 1945. On September 6 and 9 of that same year, the atomic bomb was used by the United States to destroy the Japanese cities of Hiroshima and Nagasaki.

While an atom is the smallest part of an element that still remains that element, atoms are not the smallest particles that exist. Even the protons and neutrons in the atomic nucleus are believed to be made of even smaller particles, called quarks. Current research in atomic physics focuses on describing the internal structure of atoms. By using particle accelerators, scientists are trying to characterize new particles and test the accuracy of their theories of atomic physics.

In science, a theory is not an uncertain idea or guess, but a system of related, testable ideas that satisfactorily explains some body of information. A scientific theory makes sense of facts. That all life on Earth has been produced through evolution, for example, is a fact that is explained by the theory of evolution; that all matter is made of atoms is a fact that is described by atomic theory. All scientific theories are subject to continual development and improvement, but that does not mean that there is anything fundamentally uncertain about the realities they describe.

See also Element, chemical; Nuclear fission; Nuclear fusion.

Eileen Korenic

Atomic theory

Atomic theory is the description of atoms , the smallest units of elements. The scientific evidence for the existence of atoms and its even smaller constituents is so vast that most people now consider the existence of atoms to be a fact and not just a theory.

History

Beginning in about 600 b.c., many Greek philosophers struggled to understand the nature of matter . Some said everything was made of water , which comes in three forms (solid ice , liquid water, and gaseous steam). Others believed that matter was made entirely of fire in everchanging forms. Still others believed that whatever comprised matter, it must be something that could not be destroyed but only recombined into new forms. If they could see small enough things, they would find that the same "building blocks" they started with were still there. One of these philosophers was named Democritus. He imagined starting with a large piece of matter and gradually cutting it into smaller and smaller pieces, finally reaching the smallest piece. This tiniest building block that could no longer be cut he named atomos, Greek for "no-cut." Atomos has been changed in modern times to "atom." The atoms Democritus envisioned differed only in shape and size. In his theory, different objects looked different because of the way the atoms were arranged. Aristotle, one of the most influential philosophers of that time, believed in some kind of "smallest part" of matter but not with Democritus's descriptions. Aristotle said there were only four elements (earth, air, fire, water) and that these had some smallest unit that made up all matter. Aristotle's teachings against the idea of Democritus's atom were so powerful that the idea of the atom fell out of philosophical fashion for the next 2,000 years.

Although atomic theory was abandoned for this long period, scientific experimentation, especially in chemistry , flourished. From the Middle Ages (c. 1100) onward, many chemical reactions were studied. By the seventeenth century, some of these chemists began thinking about the reactions they were seeing in terms of smallest parts. They even began using the word atom again. One of the most famous chemists of the end of the eighteenth century was Antoine Lavoisier. His chemical experiments involved very careful weighing of all the chemicals. He reacted various substances until they were in their simplest state. He found two important factors: (1) the simplest substances, which he called elements, could not be broken down any further, and (2) these elements always reacted with each other in the same proportions. The same more complex substances he called compounds. For example, two volumes of hydrogen reacted exactly with one volume of oxygen to produce water. Water could be broken down to always give exactly two volumes of hydrogen and one volume of oxygen. Lavoisier had no explanation for these amazingly consistent results. However, his numerous and careful measurements provided the clue to another chemist named John Dalton.

Dalton realized that if elements were made up of atoms, a different atom for each different element, atomic theory could explain Lavoisier's results. If two atoms of hydrogen always combined with one atom of oxygen, the resulting combination of atoms, called a molecule , would be water. Dalton published his explanation in 1803. This year is considered the beginning of modern atomic theory. Scientific experiments that followed Dalton were attempts to characterize how many elements there were, what the atoms of each element were like, how the atoms of each element were the same and how they differed, and, ultimately, whether there was anything smaller than an atom.

Describing characteristics of atoms

One of the first attributes of atoms to be described was relative atomic weight . Although a single atom was too small to weigh, atoms could be compared to each other. The chemist Jons Berzelius assumed that equal volumes of gases at the same temperature and pressure contained equal numbers of atoms. He used this idea to compare the weights of reacting gases. He was able to determine that, for example, oxygen atoms were 16 times heavier than hydrogen atoms. He made a list of these relative atomic weights for as many elements as he knew. He devised symbols for the elements by using the first letter or first two letters of their Latin names, a system still in use today. The symbol for hydrogen is H, for oxygen is O, for sodium (natrium, in Latin) is Na, and so on. The symbols also proved useful in describing how many atoms combine to form a molecule of a particular compound. For example, to show that water is made of two atoms of hydrogen and one atom of oxygen, the symbol for water is H2O. One oxygen atom can even combine with one other oxygen atom to produce a molecule of oxygen with the symbol O2 .

As more and more elements continued to be discovered, it became convenient to begin listing them in symbol form in a chart. In 1869, Dmitri Mendeleev listed the elements in order of increasing atomic weight and grouped elements that seemed to have similar chemical reactions. For example, lithium (Li), sodium (Na), and potassium (K) are all metallic elements that burst into flame if they get wet. Similar elements were placed in the same column of his chart. Mendeleev began to see a pattern among the elements, where every eighth element on the atomic weight listing would belong to the same column. Because of this periodicity or repeating pattern, Mendeleev's chart is called the "Periodic table of the elements." The table was so regular, in fact, that when there was a "hole" in the table, Mendeleev predicted that an element would eventually be discovered to fill the place. For instance, there was a space for an element with an atomic weight of about 72 (72 times heavier than hydrogen) but no known element. In 1886, 15 years after its prediction, the element Germanium (Ge) was isolated and found to have an atomic weight of 72.3. Many more elements continued to be predicted and found in this way. However, as more elements were added to the periodic table , it was found that if some elements were placed in the correct column because of similar reactions, they did not follow the right order of increasing atomic weight. Some other atomic characteristic was needed to order the elements properly. Many years passed before the correct property was found.

As chemistry experiments were searching for and characterizing more elements, other branches of science were making discoveries about electricity and light that were to contribute to the development of atomic theory. Michael Faraday had done much work to characterize electricity; James Clerk Maxwell characterized light. In the 1870s, William Crookes built an apparatus, now called a Crookes tube, to examine "rays" being given off by metals. He wanted to determine whether the rays were light or electricity based on Faraday's and Maxwell's descriptions of both. Crookes's tube consisted of a glass bulb, from which most of the air had been removed, encasing two metal plates called electrodes. One electrode was called the anode and the other was called the cathode . The plates each had a wire leading outside the bulb to a source of electricity. When electricity was applied to the electrodes, rays appeared to come from the cathode. Crookes determined that these cathode rays were particles with a negative electrical charge that were being given off by the metal of the cathode plate. In 1897, J. J. Thomson discovered that these negatively charged particles were coming out of the atoms and must have been present in the metal atoms to begin with. He called these negatively charged subatomic particles "electrons." Since the electrons were negatively charged, the rest of the atom had to be positively charged. Thomson believed that the electrons were scattered in the atom like raisins in a positively-charged bread dough, or like plums in a pudding. Although Thomson's "plum-pudding" model was not correct, it was the first attempt to show that atoms were more complex than just homogeneous spheres.

At the same time, scientists were examining other kinds of mysterious rays that were coming from the Crookes tube that did not originate at its cathode. In 1895, Wilhelm Roentgen noticed that photographic plates held near a Crookes tube would become fogged by some invisible, unknown rays. Roentgen called these rays "x rays," using "x" for unknown as in mathematics . Roentgen also established the use of photographic plates as a way to take pictures of mysterious rays. He found that by blocking the x rays with his hand, for instance, bones would block the x rays but skin and tissue would not. Doctors still use Roentgen's x rays for imaging the human body.

Photographic plates became standard equipment for scientists of Roentgen's time. One of these scientists, Henri Becquerel, left some photographic plates in a drawer with uranium , a new element he was studying. When he removed the plates, he found that they had become fogged. Since there was nothing else in the drawer, he concluded that the uranium must have been giving off some type of ray. Becquerel showed that this radiation was not as penetrating as x rays since it could be blocked by paper . The element itself was actively producing radiation, a property referred to as radioactivity. Largely through the work of Pierre and Marie Curie, more radioactive elements were found. The attempts to characterize the different types of radioactivity led to the next great chapter in the development of atomic theory.

In 1896, Ernest Rutherford, a student of J. J. Thomson, began studying radioactivity. By testing various elements and determining what kinds of materials could block the radiation from reaching a photographic plate, Rutherford concluded that there were two types of radioactivity coming from elements. He named them using the first two letters of the Greek alphabet, alpha and beta. Alpha radiation was made of positively charged particles about four times as heavy as a hydrogen atom. Beta radiation was made of negatively charged particles that seemed to be just like electrons. Rutherford decided to try an experiment using the alpha particles. He set up a piece of thin gold foil with photographic plates encircling it. He then allowed alpha particles to hit the gold. Most of the alpha particles went right through the gold foil. But a few of them did not. A few alpha particles were deflected from their straight course. A few even came straight backward. Rutherford wrote that it was as surprising as if one had fired a bullet at a piece of tissue paper only to have it bounce back. Rutherford concluded that since most of the alpha particles went through, the atoms of the gold must be mostly empty space, not Thomson's space-filling plum-pudding. Since a few of the alpha particles were deflected, there must be a densely packed positive region in each atom that he called the nucleus. With all the positive charge in the nucleus, the next question was the arrangement of the electrons in the atom.

In 1900, physicist Max Planck had been studying processes of light and heat , specifically trying to understand the light radiation given off by a "black-body," an ideal cavity made by perfectly reflecting walls. This cavity was imagined as containing objects called oscillators which absorbed and emitted light and heat. Given enough time, the radiation from such a black-body would produce a colored-light distribution called a spectrum that depended only on the temperature of the black-body and not on what it was made of. Many scientists attempted to find a mathematical relationship that would predict how the oscillators of a black-body could produce a particular spectral distribution. Max Planck found that correct mathematical relationship. He assumed that the energy absorbed or emitted by the oscillators was always a multiple of some fundamental "packet of energy" he called a quantum. Objects that emit or absorb energy do it in discrete amounts, called quanta.

At this same time, there was a physicist working with Thomson and Rutherford named Niels Bohr. Bohr realized that the idea of a quantum of energy could explain how the electrons in the atom are arranged. He described the electrons as being "in orbit" around the nucleus like planets around the sun . Like oscillators in a black-body could not have just any energy, electrons in the atom could not have just any orbit . There were only certain distances that were allowed by the energy that an electron had. If an electron of a particular atom absorbed the precisely right quantum of energy, it could move farther away from the nucleus. If an electron farther from the nucleus emitted the precisely right quantum of energy, it could move closer to the nucleus. What the precisely right values were differed for every element. These values could be determined by a process called atomic spectroscopy , an experimental technique that looked at the light spectrum produced by atoms. An atom was heated so that all of its electrons were moved far away from the nucleus. As they moved closer to the nucleus, the electrons would begin emitting their quanta of energy as light. The spectrum of light produced could be examined using a prism . The spectrum produced in this way did not show every possible color , but only those few that matched the energies corresponding to the electron orbit differences. Although later refined, Bohr's "planetary model" of the atom explained atomic spectroscopy data well enough that scientists turned their attention back to the nucleus of the atom.

Rutherford, along with Frederick Soddy, continued work with radioactive elements. Soddy, in particular, noticed that as alpha and beta particles were emitted from atoms, the atoms changed in one of two ways: (1) the element became a totally different element with completely new chemical reactions, or (2) the element maintained the same chemical reactions and the same atomic spectrum but only changed in atomic weight.

He called atoms of the second group isotopes, atoms of the same element with different atomic weights. In any natural sample of an element, there may be several types of isotopes. As a result, the atomic weight of an element that was calculated by Berzelius was actually an average of all the isotope weights for that element. This was the reason that some elements did not fall into the correct order on Mendeleev's periodic table—the average atomic weight depended on how much of each kind of isotope was present. Soddy suggested placing the elements in the periodic table by similarity of chemical reactions and then numbering them in order. The number assigned to each element in this way is called the atomic number . The atomic numbers were convenient ways to refer to elements.

Meanwhile, Thomson had continued his work with the Crookes tube. He found that, not only were cathode rays of electrons produced, but so were positive particles. After much painstaking work, he was able to separate the many different kinds of positive particles by weight. Based on these measurements, he was able to determine a fundamental particle, the smallest positive particle produced, called a proton . Since these were being produced by the atoms of the cathode and since Rutherford showed that the nucleus of the atom was positive, Thomson realized that the nucleus of an atom must contain protons. A young scientist named Henry Moseley experimented with bombarding atoms of different elements with x rays. Just as in atomic spectroscopy, where heat gives electrons more energy, x rays give protons in the nucleus more energy. And just as electrons give out light of specific energies when they cool, the nucleus emits x rays of a specific energy when it "de-excites." Moseley discovered that the energy of the emitted x rays for every element followed a simple mathematical relationship. The energy depended on the atomic number for that element, and the atomic number corresponded to the number of positive charges in the nucleus. So the correct ordering of the periodic table is by increasing number of protons in the atomic nucleus. The number of protons equals the number of electrons in a neutral atom. The electrons are responsible for the chemical reactions. Elements in the same column of the periodic table have similar arrangements of electrons with the highest energies, and this is why their reactions are similar.

Only one problem remained. Electrons had very little weight, 1/1,836 the weight of a proton. Yet the protons did not account for all of the atomic weight of an atom. It was not until 1932 that James Chadwick discovered the existence of a particle in the nucleus with no electrical charge but with a weight slightly greater than a proton. He named this particle the neutron . Neutrons are responsible for the existence of isotopes. Two atoms of the same element will have the same number of protons and electrons but they might have different numbers of neutrons and therefore different atomic weights. Isotopes are named by stating the name of the element and then the number of protons plus neutrons in the nucleus. The sum of the protons and neutrons is called the mass number . For example, uranium-235 has 235 protons and neutrons. We can look on a periodic table to find uranium's atomic number (92) which tells us the number of protons. Then by subtracting, we know that this isotope has 143 neutrons. There is another isotope of uranium, 238U, with 92 protons and 146 neutrons. Some combinations of protons and neutrons are less stable than others. Picture trying to hold 10 bowling balls in your arms. There will be some arrangement where you might be able to manage it. Now try holding 11 or only nine. There might not be a stable arrangement and you would drop the bowling balls. The same thing happens with protons and neutrons. Unstable arrangements spontaneously fall apart, emitting particles, until a stable structure is reached. This is how radioactivity like alpha particles is produced. Alpha particles are made of two protons and two neutrons tumbling out of an unstable nucleus.

Hydrogen has three kinds of isotopes: hydrogen, 2H (deuterium), and 3H (tritium).

The atomic weights of the other elements were originally compared to hydrogen without specifying which isotope. It is also difficult to get single atoms of hydrogen because it usually reacts with other atoms to form molecules like H2 or H2O. So a different element's isotope was chosen for comparison. The atomic weights are now based on 12 C (carbon-12). This isotope has six protons and six neutrons in its nucleus. Carbon-12 was defined to be 12 atomic mass units. (Atomic mass units, abbreviated amu, are units used to compare the relative weights of atoms. One amu is less than 200 sextillionths of a gram.) Every other isotope of every other element is compared to this. Then the weights of a given element's isotopes are averaged to give the atomic weights found on the periodic table.

Until this point in the story of the atom, all of the particles comprising the atom were thought of as hard, uniform spheres. Beginning in 1920 with the work of Louis de Broglie, this image changed. De Broglie showed that particles like electrons could sometimes have properties of waves. For instance, if water waves are produced by two sources, like dropping two pebbles into a pond, the waves can interfere with each other. This means that high spots add to make even higher spots. Low spots add to make even lower regions. When electrons were made to travel through a double slit, with some electrons going through one slit and some through the other, they effectively created two sources. The electrons showed this same kind of interference , producing a pattern on a collection plate. The ability of electrons and other particles to sometimes show properties of particles and sometimes of waves is called wave-particle duality. This complication to the nature of the electron meant that Bohr's idea of a planetary atom was not quite right. The electrons do have different discrete energies, but they do not follow circular orbits. In 1925, Werner Heisenberg stated that the precise speed and location of an electron cannot both be known at the same time. This "Heisenberg uncertainty principle" inspired Erwin Schrödinger to devise an equation to calculate how an electron with a certain energy moves. Schrödinger's equation describes regions in an atom where an electron with a certain energy is likely to be but not exactly where it is. This region of probability is called an orbital. Electrons move so fast within these orbitals that we can think of them as blurring into an electron cloud . Electrons move from one orbital into another by absorbing or emitting a quantum of energy, just as Bohr explained.

Applications of atomic theory

Early studies of radioactivity revealed that certain atomic nuclei were naturally radioactive. Some scientists wondered that if particles could come out of the nucleus, would it also be possible to force particles into the nucleus? In 1932, Cockcroft and Walton succeeded in building a particle accelerator, a device that could make streams of charged particles move faster and faster. These fast particles, protons for example, were then aimed at a thin plate of a lighter element like lithium (Li). If a lithium atom nucleus "captures" a proton, the nucleus becomes unstable and breaks apart into two alpha particles. This technique of inducing radioactivity by bombardment with accelerated particles is still the most used method of studying nuclear structure and subatomic particles. Today, accelerators race the particles in straight lines or, to save land space, in ringed paths several miles in diameter.

The spontaneous rearrangement of the atomic nucleus always results in a release of energy in the form of kinetic motion in fast-moving neutrons. When a large nucleus falls apart to form smaller atoms, the process is called fission. When lighter atoms are forced together to produce a heavier atom, the process is called fusion. In either case, fast neutrons are released. These can transfer their kinetic energy to the surroundings, heating it. This heat can be used to boil water, producing steam to run aturbine that turns an electric generator . Fusion is the process occurring in the center of the Sun and other stars. So much energy can be released quickly that the process has also been used for the hydrogen bomb. However, fusion is not yet controlled enough for running a power plant. Research continues to find a controlled method of using fusion energy.

On the other hand, fission reactions have also been used for very powerful weapons. The first atomic bomb was detonated in 1945. Since then, however, fission energy has also been controlled enough to operate the many nuclear power plants around the world.

While an atom is the smallest part of an element which still is that element, atoms are not the smallest particles that exist. Even the protons and neutrons in the atomic nucleus are believed to made of even smaller particles called quarks . Current research in atomic physics focuses on describing the internal structure of atoms. By using particle accelerators, scientists are trying to characterize quarks which may combine in a number of ways to produce other types of subatomic particles.

No one has ever seen a single atom even with the best optical microscopes. Special types of microscopes called scanning tunneling microscopes and atomic force microscopes make use of the forces produced by the electrons to obtain images of the electron clouds . These clouds indicate how atoms are arranged but we cannot "see" through the cloud to the nucleus. Because of the limitations of size, we will never see an atom with our own eyes. Everything we know about atoms must be deduced from larger-scale experiments. As a result, the description of atoms is still called a theory. However, this theory explains atomic experiments so well that we usually think of the existence of atoms as a fact.

See also Element, chemical; Nuclear fission; Nuclear fusion.

Eileen Korenic

KEY TERMS

Accelerator

—A device that causes particles to move faster and faster.

Alpha particle

—Two protons and two neutrons bound together and emitted from the nucleus during some kinds of radioactive decay.

Atomic mass

—The mass of an atom relative to carbon-12, 12C (which has a mass of exactly 12 atomic mass units); also the mass, in grams, of an element that contains one mole of atoms.

Atomic mass unit (u or amu)

—A unit used to express the mass of atoms equal to exactly one-twelfth of the mass of carbon-12.

Beta particle

—One type of radioactive decay particle emitted from radioactive atomic nuclei. A beta particle is the same thing as an electron.

Electrode

—A metal plate that carries electrical current.

Electron cloud

—The image of an electron moving so fast that it seems to fill a region of space.

Interference

—The combination of waves in which high spots combine to give even higher spots and low spots combine to give even lower spots.

Kinetic energy

—The energy of a moving object.

Mass number

—The sum of protons and neutrons.

Nucleus

—The dense central part of an atom containing the protons and neutrons; plural is nuclei.

Orbital

—The region of probability within an atom where an electron with a particular energy is likely to be.

Oscillators

—Objects that can absorb or emit energy and convert it into kinetic energy.

Periodicity

—Repeatability of a pattern.

Quantum

—The amount of radiant energy in the different orbits of an electron around the nucleus of an atom.

Quarks

—Believed to be the most fundamental units of protons and neutrons.

Uncertainty principle

—Heisenberg's statement that both the position and velocity of a particle cannot be known with equal precision at the same time.

Wave-particle duality

—The ability of objects to show characteristics of both waves and particles.

Atomic theory

An atomic theory is a model developed to explain the properties and behaviors of atoms. As with any scientific theory, an atomic theory is based on scientific evidence available at any given time and serves to suggest future lines of research about atoms.

The concept of an atom can be traced to debates among Greek philosophers that took place around the sixth century b.c. One of the questions that interested these thinkers was the nature of matter. Is matter, they asked, continuous or discontinuous? That is, if you could break apart a piece of chalk as long as you wanted, would you ever reach some ultimate particle beyond which further division was impossible? Or could you keep up that process of division forever? A proponent of the ultimate particle concept was the philosopher Democritus (c. 470–c. 380 b.c.), who named those particles atomos. In Greek, atomos means "indivisible."

Dalton's theory

The debate over ultimate particles was never resolved. Greek philosophers had no interest in testing their ideas with experiments. They preferred to choose those concepts that were most sound logically. For more than 2,000 years, the Democritus concept of atoms languished as kind of a secondary interest among scientists.

Then, in the first decade of the 1800s, the idea was revived. English chemist John Dalton (1766–1844) proposed the first modern atomic theory. Dalton's theory can be called modern because it contained statements about atoms that could be tested experimentally. Dalton's theory had five major parts. He said:

1. All matter is composed of very small particles called atoms.
2. All atoms of a given element are identical.
3. Atoms cannot be created, destroyed, or subdivided.
4. In chemical reactions, atoms combine with or separate from other atoms.
5. In chemical reactions, atoms combine with each other in simple, whole-number ratios to form combined atoms.

(By the term combined atoms, Dalton meant the particles that we now call molecules.)

Dalton's atomic theory is important not because everything he said was correct. It wasn't. Instead, its value lies in the research ideas it contains. As you read through the list above, you'll see that every idea can be tested by experiment.

Late nineteenth- and early twentieth-century atomic models

As each part of Dalton's theory was tested, new ideas about atoms were discovered. For example, in 1897, English physicist J. J. Thomson (1856–1940) discovered that atoms are not indivisible. When excited by means of an electrical current, atoms break down into two parts. One of those parts is a tiny particle carrying a negative electrical charge, the electron.

To explain what he had discovered, Thomson suggested a new model of the atom, a model widely known as the plum-pudding atom. The name comes from a comparison of the atom with a traditional English plum pudding, in which plums are embedded in pudding, as shown in the accompanying figure of the evolution of atomic theory. In Thomson's atomic model, the "plums" are negatively charged electrons, and the "pudding" is a mass of positive charge.

The nuclear atom. Like the Dalton model before it, Thomson's plumpudding atom was soon put to the test. It did not survive very long. In the period between 1906 and 1908, English chemist and physicist Ernest Rutherford (1871–1937) studied the effects of bombarding thin gold foil with alpha particles. Alpha particles are helium atoms that have lost their electrons and that, therefore, are positively charged. Rutherford reasoned that the way alpha particles traveled through the gold foil would give him information about the structure of gold atoms in the foil.

Rutherford's experiments provided him with two important pieces of information. First, most of the alpha particles traveled right through the foil without being deflected at all. This result tells us, Rutherford concluded, that atoms consist mostly of empty space. Second, a few of the alpha particles were deflected at very sharp angles. In fact, some reflected completely backwards and were detected next to the gun from which they were first produced. Rutherford was enormously surprised. The result, he said, was something like shooting a cannon ball at a piece of tissue paper and having the ball bounce back at you.

According to Rutherford, the conclusion to be drawn from this result was that the positive charge in an atom must all be packed together in one small region of the atom. He called that region the nucleus of the atom. A sketch of Rutherford's nuclear atom is shown in the figure as well.

The planetary atom. One part of Rutherford's model—the nucleus—has turned out to be correct. However, his placement of electrons created some problems, which he himself recognized. The peculiar difficulty is that electrons cannot remain stationary in an atom, as they appear to be in the figure. If they were stationary, they would be attracted to the nucleus and become part of it. (Remember that electrons are negatively charged and the nucleus is positively charged; opposite charges attract.)

But the electrons could not be spinning around the nucleus either. According to a well-known law of physics, charged particles (like electrons) that travel through space give off energy. Moving electrons would eventually lose energy, lose speed, and fall into the nucleus. Electrons in Rutherford's atom could neither be at rest nor in motion.

The solution to this dilemma was proposed in a new and brilliant atomic theory in 1913. Suppose, said Danish physicist Niels Bohr (1885–1962), that places exist in the atom where electrons can travel without losing energy. Let's call those places "permitted orbits," something like the orbits that planets travel in their journey around the Sun. A sketch of Bohr's planetary atom is also shown in the figure. If we can accept that idea, Bohr said, the problem with electrons in Rutherford's atom would be solved.

Scientists were flabbergasted. Bohr was saying that the way to explain the structure of an atom was to ignore an accepted principle of physics—at least for certain small parts of the atom. The Bohr model sounded almost like cheating: inventing a model just because it might look right.

The test, of course, was to see if the Bohr model could survive experiments designed specifically to test it. And it did. Within a very short period of time, other scientists were able to report that the Bohr model met all the tests they were able to devise for it. By 1930, then, the accepted model of the atom consisted of two parts, a nucleus whose positive charge was known to be due to tiny particles called protons, and one or more electrons arranged in distinct orbits outside the nucleus.

The neutron. One final problem remained. In the Bohr model, there must be an equal number of protons and electrons. This balance is the only way to be sure that an atom is electrically neutral, which we know to be the case for all atoms. But if one adds up the mass (total amount of matter) of all the protons and electrons in an atom, the total comes no where near the actual mass of an atom.

The solution to this problem was suggested by English physicist James Chadwick (1891–1974) in 1932. The reason for mass differences, Chadwick found, was that the nuclei of atoms contain a particle with no electric charge. He called this particle a neutron.

Chadwick's discovery resulted in a model of the atom that is fairly easy to understand. The core of the atom is the atomic nucleus, in which are found one or more protons and neutrons. Outside the nucleus are electrons traveling in discrete orbits.

Modern theories

This model of the atom can be used to explain many of the ideas in chemistry in which ordinary people are interested. But the model has not been used by chemists themselves for many decades. The reason for this difference is that revolutionary changes occurred in physics during the 1920s. These changes included the rise of relativity, quantum theory, and uncertainty that forced chemists to rethink the most basic concepts about atoms.

As an example, the principle of uncertainty says that it is impossible to describe with perfect accuracy both the position and the motion of an object. In other words, you might be able to say very accurately where an electron is located in an atom, but to do so reduces the accuracy with which you can describe its motion.

By the end of the 1920s, then, chemists had begun to look for new ways to describe the atom that would incorporate the new discoveries in physics. One step in this direction was to rely less on physical models and more on mathematical models. That is, chemists began to give up on the idea of an electron as a tiny particle carrying an electrical charge traveling in a certain direction with a certain speed in a certain part of an atom. Instead, they began to look for mathematical equations which, when solved, gave the correct answers for the charge, mass, speed, spin, and other properties of the electron.

Mathematical models of the atom are often very difficult to understand, but they are enormously useful and successful for professional chemists. The clues they have given about the ultimate structure of matter have led not only to a better understanding of atoms themselves, but also to the development of countless innovative new products in our daily lives.

Do atoms exist?

One of the most remarkable features of atomic theory is that even today, after hundreds of years of research, no one has yet seen a single atom. Some of the very best microscopes have produced images of groups of atoms, but no actual picture of an atom yet exists. How, then, can scientists be so completely certain of the existence of atoms and of the models they have created for them? The answer is that models of the atom, like other scientific models, can be tested by experimentation. Those models that pass the test of experimentation survive, while those that do not are abandoned. The model of atoms that scientists use today has survived and been modified by untold numbers of experiments and will be subjected to other such tests in the future.

[See also Atom; Atomic mass; Electron; Element, chemical; Isotope; Periodic table; Subatomic particle ]

Atomic theory

One of the points of dispute among early Greek philosophers was the ultimate nature of matter. The question was whether the characteristics of matter that can be observed with the five senses are a true representation of matter at its most basic level. Some philosophers thought that they were. Anaxagoras of Klazomenai (c. 498–428 b.c.), for example, taught that matter can be sub-divided without limit and that it retains its characteristics no matter how it is divided.

An alternative view was that of Leucippus of Miletus (about 490 b.c.) and his pupil, Democritus of Abdera (c. 460–370 b.c.). The views of these scholars are preserved in a few fragments of their writings and of commentaries on their teachings. Some writers doubt that Leucippus even existed. In any case, the ideas attributed to them are widely known. They thought that all matter consists of tiny, indivisible particles moving randomly about in a void (a vacuum). The particles were described as hard, with form and size, but no color, taste, or smell. They became known by the Greek word atomos, meaning "indivisible." Democritus suggested that, from time to time, atoms collide and combine with each other by means of hook-and-eye attachments on their surfaces.

Perhaps the most effective popularizer of the atomic theory was the Roman poet and naturalist, Lucretius. In his poem De Rerum Natura (On the nature of things) Lucretius states that only two realities exist, solid, everlasting particles and the void. This atomistic philosophy was in competition with other ideas about the fundamental nature of matter. Aristotle, for example, rejected Democritus' ideas because he could not accept the concept of a vacuum nor the idea that particles could move about on their own.

In addition, debates between atomists and anti-atomists quickly developed religious overtones. As the natural philosophy of Aristotle was adopted by and incorporated into early Christian theology, anti-atomism became acceptable and "correct," atomism, heretical. In fact, one objective of Lucretius' poem was to provide a materialistic explanation of the world designed to counteract religious superstition rampant at the time.

In spite of official disapproval, the idea of fundamental particles held a strong appeal for at least some philosophers through the ages. The French philosopher Pierre Gassendi (1592–1655) was especially influential in reviving and promoting the idea of atomism. Robert Boyle and Isaac Newton were both enthusiastic supporters of the theory.

Credit for the first modern atomic theory goes to the English chemist, John Dalton. In his 1808 book, A New System of Chemical Philosophy, Dalton outlined five fundamental postulates about atoms: 1. All matter consists of tiny, indivisible particles, which Dalton called atoms. 2. All atoms of a particular element are exactly alike, but atoms of different elements are different. 3. All atoms are unchangeable. 4. Atoms of elements combine to form "compound atoms" (i.e., molecules) of compounds. 5. In chemical reactions, atoms are neither created nor destroyed, but are only rearranged.

A key distinguishing feature of Dalton's theory was his emphasis on the weights of atoms. He argued that every atom had a specific weight that could be determined by experimental analysis. Although the specific details of Dalton's proposed mechanism for determining atomic weights were flawed, his proposal stimulated other chemists to begin research on atomic weights.

Dalton's theory was widely accepted because it explained so many existing experimental observations and because it was so fruitful in suggesting new lines of research. But the theory proved to be wrong in many of its particulars. For example, in 1897, the English physicist Joseph J. Thomson showed that particles even smaller than the atom, electrons, could be extracted from atoms. Atoms could not, therefore, be indivisible. The discovery of radioactivity at about the same time showed that at least some atoms are not unchangeable, but instead, spontaneously decay into other kinds of atoms.

By 1913, the main features of the modern atomic theory had been worked out. The work of Ernest Rutherford, Niels Bohr and others, suggested that an atom consists of a central core, the nucleus, surrounded by one or more electrons, arranged in energy levels each of which can hold some specific number of electrons.

Bohr's atomic model marked the beginning of a new approach in constructing atomic theory. His work, along with that of Erwin Schrödinger , Louis Victor de Broglie, Werner Karl Heisenberg, Paul Adrien Maurice Dirac, and others showed that atoms could be understood and represented better through mathematics than through physical models. Instead of drawing pictures that show the location and movement of particles within the atoms, modern scientists tend to write mathematical equations that describe the behavior of observed atomic phenomena.

See also Atomic mass and weight; Atomic number; Chemical bonds and physical properties

atomic theory the theory that all matter is made up of tiny indivisible particles or atoms, which in ancient times was taught notably by Democritus and Epicurus. According to the modern version, the atoms of each element are effectively identical, but differ from those of other elements, and unite to form compounds in fixed proportions.