Our world is made up of atoms, yet the atomic model of the universe is nonetheless considered a "theory." When scientists know beyond all reasonable doubt that a particular principle is the case, then it is dubbed a law. Laws address the fact that certain things happen, as well as how they happen. A theory, on the other hand, attempts to explain why things happen. By definition, an idea that is dubbed a theory has yet to be fully proven, and such is the case with the atomic theory of matter. After all, the atom cannot be seen, even with electron microscopes—yet its behavior can be studied in terms of its effects. Atomic theory explains a great deal about the universe, including the relationship between chemical elements, and therefore (as with Darwin's theory concerning biological evolution), it is generally accepted as fact. The particulars of this theory, including the means by which it evolved over the centuries, are as dramatic as any detective story. Nonetheless, much still remains to be explained about the atom—particularly with regard to the smallest items it contains.
HOW IT WORKS
Why Study Atoms?
Many accounts of the atom begin with a history of the growth in scientists' understanding of its structure, but here we will take the opposite approach, first discussing the atom in terms of what physicists and chemists today understand. Only then will we examine the many challenges scientists faced in developing the current atomic model: false starts, wrong theories, right roads not taken, incomplete models. In addition, we will explore the many insights added along the way as, piece by piece, the evidence concerning atomic behavior began to accumulate.
People who are not scientifically trained tend to associate studies of the atom with physics, not chemistry. While it is true that physicists study atomic structure, and that much of what scientists know today about atoms comes from the work of physicists, atomic studies are even more integral to chemistry than to physics. At heart, chemistry is about the interaction of different atomic and molecular structures: their properties, their reactions, and the ways in which they bond.
WHAT THE ATOM MEANS TO CHEMISTRY.
Just as a writer in English works with the 26 letters of the alphabet, a chemist works with the 100-plus known elements, the fundamental and indivisible substances of all matter. And what differentiates the elements, ultimately, from one another is not their color or texture, or even the phase of matter—solid, gas, or liquid—in which they are normally found. Rather, the defining characteristic of an element is the atom that forms its basic structure.
The number of protons in an atom is the critical factor in differentiating between elements, while the number of neutrons alongside the protons in the nucleus serves to distinguish one isotope from another. However, as important as elements and even isotopes are to the work of a chemist, the components of the atom's nucleus have little direct bearing on the atomic activity that brings about chemical reactions and chemical bonding. All the chemical "work" of an atom is done by particles vastly smaller in mass than either the protons or neutrons—fast-moving little bundles of energy called electrons.
Moving rapidly through the space between the nucleus and the edge of the atom, electrons sometimes become dislodged, causing the atom to become a positively charged ion. Conversely, sometimes an atom takes on one or more electrons, thus acquiring a negative charge. Ions are critical to the formation of some kinds of chemical bonds, but the chemical role of the electron is not limited to ionic bonds.
In fact, what defines an atom's ability to bond with another atom, and therefore to form a molecule, is the specific configuration of its electrons. Furthermore, chemical reactions are the result of changes in the arrangement of electrons, not of any activity involving protons or neutrons. So important are electrons to the interactions studied in chemistry that a separate essay has been devoted to them.
What an Atom Is
BASIC ATOMIC STRUCTURE.
The definitions of atoms and elements seems, at first glance, almost circular: an element is a substance made up of only one kind of atom, and an atom is the smallest particle of an element that retains all the chemical and physical properties of the element. In fact, these two definitions do not form a closed loop, as they would if it were stated that an element is something made up of atoms. Every item of matter that exists, except for the subatomic particles discussed in this essay, is made up of atoms. An element, on the other hand, is—as stated in its definition—made up of only one kind of atom. "Kind of atom" in this context refers to the number of protons in its nucleus.
Protons are one of three basic subatomic particles, the other two being electrons and neutrons. As we shall see, there appear to be particles even smaller than these, but before approaching these "sub-subatomic" particles, it is necessary to address the three most significant components of an atom. These are distinguished from one another in terms of electric charge: protons are positively charged, electrons are negative in charge, and neutrons have no electrical charge. As with the north and south poles of magnets, positive and negative charges attract one another, whereas like charges repel. Atoms have no net charge, meaning that the protons and electrons cancel out one another.
EVOLVING MODELS OF THE ATOM.
Scientists originally thought of an atom as a sort of closed sphere with a relatively hard shell, rather like a ball bearing. Nor did they initially understand that atoms themselves are divisible, consisting of the parts named above. Even as awareness of these three parts emerged in the last years of the nineteenth century and the first part of the twentieth, it was not at all clear how they fit together.
At one point, scientists believed that electrons floated in a cloud of positive charges. This was before the discovery of the nucleus, where the protons and neutrons reside at the heart of the atom. It then became clear that electrons were moving around the nucleus, but how? For a time, a planetary model seemed appropriate: in other words, electrons revolved around the nucleus much as planets orbit the Sun. Eventually, however—as is often the case with scientific discovery—this model became unworkable, and had to be replaced by another.
The model of electron behavior accepted today depicts the electrons as forming a cloud around the nucleus—almost exactly the opposite of what physicists believed a century ago. The use of the term "cloud" may perhaps be a bit misleading, implying as it does something that simply hovers. In fact, the electron, under normal circumstances, is constantly moving. The paths of its movement around the nucleus are nothing like that of a planet's orbit, except inasmuch as both models describe a relatively small object moving around a relatively large one.
The furthest edges of the electron's movement define the outer perimeters of the atom. Rather than being a hard-shelled little nugget of matter, an atom—to restate the metaphor mentioned above—is a cloud of electrons surrounding a nucleus. Its perimeters are thus not sharply delineated, just as there is no distinct barrier between Earth's atmosphere and space itself. Just as the air gets thinner the higher one goes, so it is with an atom: the further a point is from the nucleus, the less the likelihood that an electron will pass that point on a given orbital path.
MASS NUMBER AND ATOMIC NUMBER.
The term nucleon is used generically to describe the relatively heavy particles that make up an atomic nucleus. Just as "sport" can refer to football, basketball, or baseball, or any other item in a similar class, such as soccer or tennis, "nucleon" refers to protons and neutrons. The sum of protons and neutrons is sometimes called the nucleon number, although a more commonly used term is mass number.
Though the electron is the agent of chemical reactions and bonding, it is the number of protons in the nucleus that defines an atom as to its element. Atoms of the same element always have the same number of protons, and since this figure is unique for a given element, each element is assigned an atomic number equal to the number of protons in its nucleus. The atoms are listed in order of atomic number on the periodic table of elements.
ATOMIC MASS AND ISOTOPES.
A proton has a mass of 1.673 · 10−24 g, which is very close to the established figure for measuring atomic mass, the atomic mass unit. At one time, the basic unit of atomic mass was equal to the mass of one hydrogen atom, but hydrogen is so reactive—that is, it tends to combine readily with other atoms to form a molecule, and hence a compound—that it is difficult to isolate. Instead, the atomic mass unit is today defined as 1/12 of the mass of a carbon-12 atom. That figure is exactly 1.66053873 · 10−24 grams.
The mention of carbon-12, a substance found in all living things, brings up the subject of isotopes. The "12" in carbon-12 refers to its mass number, or the sum of protons and neutrons. Two atoms may be of the same element, and thus have the same number of protons, yet differ in their number of neutrons—which means a difference both in mass number and atomic mass. Such differing atoms of the same element are called isotopes. Isotopes are often designated by symbols showing mass number to the upper left of the chemical symbol or element symbol—for instance, 12C for carbon-12.
Protons have a positive electric charge of 1, designated either as 1+ or +1. Neutrons, on the other hand, have no electric charge. It appears that the 1+ charge of a proton and the 0 charge of a neutron are the products of electric charges on the part of even smaller particles called quarks. A quark may either have a positive electric charge of less than 1+, in which case it is called an "up quark"; or a negative charge of less than 1−, in which case it is called a "down quark."
Research indicates that a proton contains two up quarks, each with a charge of 2/3+, and one down quark with a charge of 1/3−. This results in a net charge of 1+. On the other hand, a neutron is believed to hold one up quark with a charge of 2/3+, and two down quarks with charges of 1/3− each. Thus, in the neutron, the up and down quarks cancel out one another, and the net charge is zero.
A neutron has about the same mass as a proton, but other than its role in forming isotopes, the neutron's function is not exactly clear. Perhaps, it has been speculated, it binds protons—which, due to their positive charges, tend to repel one another—together at the nucleus.
An electron is much smaller than a proton or neutron, and has much less mass; in fact, its mass is equal to 1/1836 that of a proton, and 1/1839 that of a neutron. Yet the area occupied by electrons—the region through which they move—constitutes most of the atom's volume. If the nucleus of an atom were the size of a BB (which, in fact, is billions of times larger than a nucleus), the furthest edge of the atom would be equivalent to the highest ring of seats around an indoor sports arena. Imagine the electrons as incredibly fast-moving insects buzzing constantly through the arena, passing by the BB but then flitting to the edges or points in between, and you have something approaching an image of the atom's interior.
How fast does an electron move? Speeds vary depending on a number of factors, but it can move nearly as fast as light: 186,000 mi (299,339 km) per second. On the other hand, for an item of matter near absolute zero in temperature, the velocity of the electron is much, much less. In any case, given the fact that an electron has enough negative charge to cancel out that of the proton, it must be highly energized. After all, this would be like an electric generator weighing 1 lb having as much power as a generator that weighed 1 ton.
According to what modern scientists know or hypothesize concerning the inner structure of the atom, electrons are not made up of quarks; rather, they are part of a class of particles called leptons. It appears that leptons, along with quarks and what are called exchange particles, constitute the elementary particles of atoms—particles on a much more fundamental level than that of the proton and neutron.
Electrons are perhaps the most intriguing parts of an atom. Their mass is tiny, even in atomic terms, yet they possess enough charge to counteract a "huge" proton. They are capable, in certain situations, of moving from one atom to another, thus creating ions, and depending on their highly complex configuration and ability to rearrange their configuration, they facilitate or prevent chemical reactions.
Ancient Greek Theories of Matter
The first of the Greek philosophers, and the first individual in Western history who deserves to be called a scientist, was Thales (c. 625-c. 547 b.c.) of Miletus. (Miletus is in Greek Asia Minor, now part of Turkey.) Among his many achievements were the correct prediction of a solar eclipse, and one of the first-ever observations of electricity, when he noted the electrification of amber by friction.
But perhaps the greatest of Thales's legacies was his statement that "Everything is water." This represented the first attempt to characterize the nature of all physical reality. It set off a debate concerning the fundamental nature of matter that consumed Greek philosophers for two centuries. Later, philosophers attempted to characterize matter in terms of fire or air. In time, however, there emerged a school of thought concerned not with identifying matter as one particular thing or another, but with recognizing a structural consistency in all of matter. Among these were the philosophers Leucippus (c. 480-c. 420 b.c.) and his student Democritus (c. 460-370 b.c.)
Leucippus and Democritus proposed a new and highly advanced model for the tiniest point of physical space. Democritus, who actually articulated these ideas (far less is known about Leucippus) began with a "thought experiment," imagining what would happen if an item of matter were subdivided down to its smallest piece. This tiniest fragment, representing an item of matter that could not be cut into smaller pieces, he called by a Greek term meaning "no cut": atomos.
Democritus was not necessarily describing matter in a concrete, scientific way: his "atoms" were idealized philosophical constructs rather than purely physical units. Yet, he came amazingly close, and indeed much closer than any thinker for the next 22 centuries, to identifying the fundamental structure of physical reality. Why did it take so long for scientists to come back around to the atomic model? The principal culprit, who advanced an erroneous theory of matter, also happened to be one of the greatest thinkers of all time: Aristotle (384-322 b.c..)
Aristotle made numerous contributions to science, including his studies in botany and zoology, as well as his explanation of the four causes, a significant attempt to explain events by means other than myth or superstition. In the area of the physical sciences, however, Aristotle's impact was less than beneficial. Most notably, in explaining why objects fall when dropped, he claimed that the ground was their "natural" destination—a fallacy later overturned with the gravitational model developed by Galileo Galilei (1564-1642) and Sir Isaac Newton (1642-1727).
The ideas Aristotle put forward concerning what he called "natural motion" were a product of his equally faulty theories with regard to what today's scientists refer to as chemistry. In ancient times, chemistry, as such, did not exist. Long before Aristotle's time, Egyptian embalmers and metallurgists used chemical processes, but they did so in a practical, applied manner, exerting little effort toward what could be described as scientific theory. Philosophers such as Aristotle, who were some of the first scientists, made little distinction between physical and chemical processes. Thus, whereas physics is understood today as an important background for chemistry, Aristotle's "physics" was actually an outgrowth of his "chemistry."
Rejecting Democritus's atomic model, Aristotle put forward his own view of matter. Like Democritus, he believed that matter was composed of very small components, but these he identified not as atoms, but as "elements": earth, air, fire, and water. He maintained that all objects consisted, in varying degrees, of one or more of these, and based his explanation of gravity on the relative weights of each element. Water sits on top of the earth, he explained, because it is lighter, yet air floats above the water because it is lighter still—and fire, lightest of all, rises highest. Furthermore, he claimed that the planets beyond Earth were made up of a "fifth element," or quintessence, of which little could be known.
In fairness to Aristotle, it should be pointed out that it was not his fault that science all but died out in the Western world during the period from about a.d. 200 to about 1200. Furthermore, he did offer an accurate definition of an element, in a general sense, as "one of those simple bodies into which other bodies can be decomposed, and which itself is not capable of being divided into others." As we shall see, the definition used today is not very different from Aristotle's. However, to define an element scientifically, as modern chemists do, it is necessary to refer to something Aristotle rejected: the atom. So great was his opposition to Democritus's atomic theory, and so enormous was Aristotle's influence on learning for more than 1,500 years following his death, that scientists only began to reconsider atomic theory in the late eighteenth century.
A Maturing Concept of Elements
BOYLE'S IDEA OF ELEMENTS.
One of the first steps toward an understanding of the chemical elements came with the work of English physicist and chemist Robert Boyle (1627-1691). Building on the usable definition of an element provided by Aristotle, Boyle maintained no substance was an element if it could be broken down into other substances. Thus, air could be eliminated from the list of "elements," because, clearly, it could be separated into more than one elemental substance. (In fact, none of the four "elements" identified by Aristotle even remotely qualifies as an element in modern chemistry.)
Boyle, nonetheless, still clung to aspects of alchemy, a pseudo-science based on the transformation of "base metals," for example, the metamorphosis of iron into gold. Though true chemistry grew out of alchemy, the fundamental proposition of alchemy was faulty: if one metal can be turned into another, then that means that metals are not elements, which, in fact, they are. Nonetheless, Boyle's studies led to the identification of numerous elements—that is, items that really are elements—in the years that followed.
LAVOISIER AND PROUST: CONSTANT COMPOSITION.
A few years after Boyle came two French chemists who extended scientific understanding of the elements. Antoine Lavoisier (1743-1794) affirmed the definition of an element as a simple substance that could not be broken down into a simpler substance, and noted that elements always react with one another in the same proportions.
Joseph-Louis Proust (1754-1826) put forward the law of constant composition, which holds that a given compound always contains the same proportions of mass between elements. Another chemist of the era had claimed that the composition of a compound varies in accordance with the reactants used to produce it. Proust's law of constant composition made it clear that any particular compound will always have the same composition.
Early Modern Understanding of the Atom
DALTON AND AVOGADRO: ATOMS AND MOLECULES.
The work of Lavoisier and Proust influenced a critical figure in the development of the atomic model: English chemist John Dalton (1766-1844). In A New System of Chemical Philosophy (1808), Dalton put forward the idea that nature is composed of tiny particles, and in so doing he adopted Democritus's word "atom" to describe these basic units. This idea, which Dalton had formulated five years earlier, marked the starting-point of modern atomic theory.
Dalton recognized that the structure of atoms in a particular element or compound is uniform, but maintained that compounds are made up of compound atoms: in other words, water, for instance, is a compound of "water atoms." However, water is not an element, and thus, it was necessary to think of its atomic composition in a different way—in terms of molecules rather than atoms. Dalton's contemporary Amedeo Avogadro (1776-1856), an Italian physicist, became the first scientist to clarify the distinction between atoms and molecules.
The later development of the mole, which provided a means whereby equal numbers of molecules could be compared, paid tribute to Avogadro by designating the number of molecules in a mole as "Avogadro's number." Another contemporary, Swedish chemist Jons Berzelius (1779-1848), maintained that equal volumes of gases at the same temperature and pressure contained equal numbers of atoms. Using this idea, he compared the mass of various reacting gases, and developed a system of comparing the mass of various atoms in relation to the lightest one, hydrogen. Berzelius also introduced the system of chemical symbols—H for hydrogen, O for oxygen, and so on—in use today.
BROWNIAN MOTION AND KINETIC THEORY.
Yet another figure whose dates overlapped with those of Dalton, Avogadro, and Berzelius was Scottish botanist Robert Brown (1773-1858). In 1827, Brown noted a phenomenon that later had an enormous impact on the understanding of the atom. While studying pollen grains under a microscope, Brown noticed that the grains underwent a curious zigzagging motion in the water. The pollen assumed the shape of a colloid, a pattern that occurs when particles of one substance are dispersed—but not dissolved—in another substance. At first, Brown assumed that the motion had a biological explanation—that is, it resulted from life processes within the pollen—but later, he discovered that even pollen from long-dead plants behaved in the same way.
Brown never understood what he was witnessing. Nor did a number of other scientists, who began noticing other examples of what came to be known as Brownian motion: the constant but irregular zigzagging of colloidal particles, which can be seen clearly through a microscope. Later, however, Scottish physicist James Clerk Maxwell (1831-1879) and others were able to explain this phenomenon by what came to be known as the kinetic theory of matter.
Kinetic theory is based on the idea that molecules are constantly in motion: hence, the water molecules were moving the pollen grains Brown observed. Pollen grains are many thousands of times as large as water molecules, but since there are so many molecules in even a drop of water, and their motion is so constant but apparently random, they are bound to move a pollen grain once every few thousand collisions.
Mendeleev and the Periodic Table
In 1869, Russian chemist Dmitri Mendeleev (1834-1907) introduced a highly useful system for organizing the elements, the periodic table. Mendeleev's table is far more than just a handy chart listing elements: at once simple and highly complex, it shows elements in order of increasing atomic mass, and groups together those exhibiting similar forms of chemical behavior and structure.
Reading from right to left and top to bottom, the periodic table, as it is configured today, lists atoms in order of atomic number, generally reflected by a corresponding increase in average atomic mass. As Mendeleev observed, every eighth element on the chart exhibits similar characteristics, and thus the chart is organized in columns representing specific groups of elements.
The patterns Mendeleev observed were so regular that for any "hole" in his table, he predicted that an element would be discovered that would fill that space. For instance, at one point there was a gap between atomic numbers 71 and 73 (lutetium and tantalum, respectively). Mendeleev indicated that an atom would be found for the space, and 15 years after this prediction, the element germanium was isolated.
However, much of what defines an element's place on the chart today relates to subatomic particles—protons, which determine atomic number, and electrons, whose configurations explain certain chemical similarities. Mendeleev was unaware of these particles: from the time he created his table, it was another three decades before the discovery of the first of these particles, the electron. Instead, he listed the elements in an order reflecting outward characteristics now understood to be the result of the quantity and distribution of protons and electrons.
Electromagnetism and Radiation
The contribution of Mendeleev's contemporary, Maxwell, to the understanding of the atom was not limited to his kinetic theory. Building on the work of British physicist and chemist Michael Faraday (1791-1867) and others, in 1865 he published a paper outlining a theory of a fundamental interaction between electricity and magnetism. The electromagnetic interaction, as it later turned out, explained something that gravitation, the only other form of fundamental interaction known at the time, could not: the force that held together particles in an atom.
The idea of subatomic particles was still a long time in coming, but the model of electromagnetism helped make it possible. In the long run, electromagnetism was understood to encompass a whole spectrum of energy radiation, including radio waves; infrared, visible, and ultraviolet light; x rays; and gamma rays. But this, too, was the product of work on the part of numerous individuals, among whom was English physicist William Crookes (1832-1919).
In the 1870s, Crookes developed an apparatus later termed a Crookes tube, with which he sought to analyze the "rays"—that is, radiation—emitted by metals. The tube consisted of a glass bulb, from which most of the air had been removed, encased between two metal plates or electrodes, referred to as a cathode and an anode. A wire led outside the bulb to an electric source, and when electricity was applied to the electrodes, the cathodes emitted rays. Crookes concluded that the cathode rays were particles with a negative electric charge that came from the metal in the cathode plate.
In 1895, German physicist Wilhelm Röntgen (1845-1923) noticed that photographic plates held near a Crookes tube became fogged, and dubbed the rays that had caused the fogging "x rays." A year after Röntgen's discovery, French physicist Henri Becquerel (1852-1908) left some photographic plates in a drawer with a sample of uranium. Uranium had been discovered more than a century before; however, there were few uses for it until Becquerel discovered that the uranium likewise caused a fogging of the photographic plates.
Thus radioactivity, a type of radiation brought about by atoms that experience radioactive decay was discovered. The term was coined by Polish-French physicist and chemist Marie Curie (1867-1934), who with her husband Pierre (1859-1906), a French physicist, was responsible for the discovery of several radioactive elements.
The Rise and Fall of the Plum Pudding Model
Working with a Crookes tube, English physicist J. J. Thomson (1856-1940) hypothesized that the negatively charged particles Crookes had observed were being emitted by atoms, and in 1897, he gave a name to these particles: electrons. The discovery of the electron raised a new question: if Thomson's particles exerted a negative charge, from whence did the counterbalancing positive charge come?
An answer, of sorts, came from William Thomson, not related to the other Thomson and, in any case, better known by his title as Lord Kelvin (1824-1907). Kelvin compared the structure of an atom to an English plum pudding: the electrons were like raisins, floating in a positively charged "pudding"—that is, an undifferentiated cloud of positive charges.
Kelvin's temperature scale contributed greatly to the understanding of molecular motion as encompassed in the kinetic theory of matter. However, his model for the distribution of charges in an atom—charming as it may have been—was incorrect. Nonetheless, for several decades, the "plum pudding model," as it came to be known, remained the most widely accepted depiction of the way that electric charges were distributed in an atom. The overturning of the plum pudding model was the work of English physicist Ernest Rutherford (1871-1937), a student of J. J. Thomson.
RUTHERFORD IDENTIFIES THE NUCLEUS.
Rutherford did not set out to disprove the plum pudding model; rather, he was conducting tests to find materials that would block radiation from reaching a photographic plate. The two materials he identified, which were, respectively, positive and negative in electric charge, he dubbed alpha and beta particles. (An alpha particle is a helium nucleus stripped of its electrons, such that it has a positive charge of 2; beta particles are either electrons or positively charged subatomic particles called positrons. The beta particle Rutherford studied was an electron emitted during radioactive decay.)
Using a piece of thin gold foil with photographic plates encircling it, Rutherford bombarded the foil with alpha particles. Most of the alpha particles went straight through the foil—as they should, according to the plum pudding model. However, a few particles were deflected from their course, and some even bounced back. Rutherford later said it was as though he had fired a gun at a piece of tissue paper, only to see the tissue deflect the bullets. Analyzing these results, Rutherford concluded that there was no "pudding" of positive charges: instead, the atom had a positively charged nucleus at its center.
The Nucleus Emerges
PROTONS AND ISOTOPES.
In addition to defining the nucleus, Rutherford also gave a name to the particles that imparted its positive charge: protons. But just as the identification of the electron had raised new questions that, in being answered, led to the discovery of the proton, Rutherford's achievement only brought up new anomalies concerning the behavior of the nucleus.
Together with English chemist Frederick Soddy (1877-1956), Rutherford discovered that when an atom emitted alpha or beta particles, its atomic mass changed. Soddy had a name for atoms that displayed this type of behavior: isotopes. Certain types of isotopes, Soddy and Rutherford went on to conclude, had a tendency to decay, moving toward stabilization, and this decay explained radioactivity.
CLARIFYING THE PERIODIC TABLE.
Soddy concluded that atomic mass, as measured by Berzelius, was actually an average of the mass figures for all isotopes within that element. This explained a problem with Mendeleev's periodic table, in which there seemed to be irregularities in the increase of atomic mass from element to element. The answer to these variations in mass, it turned out, related to the number of isotopes associated with a given element: the greater the number of isotopes, the more these affected the overall measure of the element's mass.
By this point, physicists and chemists had come to understand that various levels of energy in matter emitted specific electromagnetic wavelengths. Welsh physicist Henry Moseley (1887-1915) experimented with x rays, bombarding atoms of different elements with high levels of energy and observing the light they gave off as they cooled. In the course of these tests, he uncovered an astounding mathematical relationship: the amount of energy a given element emitted was related to its atomic number.
Furthermore, the atomic number corresponded to the number of positive charges—this was in 1913, before Rutherford had named the proton—in the nucleus. Mendeleev had been able to predict the discovery of new elements, but such predictions had remained problematic. When scientists understood the idea of atomic number, however, it became possible to predict the existence of undiscovered elements with much greater accuracy.
Yet again, discoveries—the nucleus, protons, and the relationship between these and atomic number—only created new questions. (This, indeed, is one of the hallmarks of an active scientific theory. Rather than settling questions, science is about raising new ones, and thus improving the quality of the questions that are asked.) Once Rutherford had identified the proton, and Moseley had established the number of protons, the mystery at the heart of the atom only grew deeper.
Scientists had found that the measured mass of atoms could not be accounted for by the number of protons they contained. Certainly, the electrons had little to do with atomic mass: by then it had been shown that the electron weighed about 0.06% as much as a proton. Yet for all elements other than protium (the first of three hydrogen isotopes), there was a discrepancy between atomic mass and atomic number. Clearly, there had to be something else inside the nucleus.
In 1932, English physicist James Chadwick (1891-1974) identified that "something else." Working with radioactive material, he found that a certain type of subatomic particle could penetrate lead. All other known types of radiation were stopped by the lead, and therefore, Chadwick reasoned that this particle must be neutral in charge. In 1932, he won the Nobel Prize in Physics for his discovery of the neutron.
The Nuclear Explosion
ISOTOPES AND RADIOACTIVITY.
Chadwick's discovery clarified another mystery, that of the isotope, which had been raised by Rutherford and Soddy several decades earlier. Obviously, the number of protons in a nucleus did not change, but until the identification of the neutron, it had not been clear what it was that did change. At that point, it was understood that two atoms may have the same atomic number—and hence be of the same element—yet they may differ in number of neutrons, and thus be isotopes.
As the image of what an isotope was became clearer, so too did scientists' comprehension of radioactivity. Radioactivity, it was discovered, was most intense where an isotope was the most unstable—that is, in cases where an isotope had the greatest tendency to experience decay. Uranium had a number of radioactive isotopes, such as
235U, and these found application in the burgeoning realm of nuclear power—both the destructive power of atomic bombs, and later the constructive power of nuclear energy plants.
FISSION VS. FUSION.
In nuclear fission, or the splitting of atoms, uranium isotopes (or other radioactive isotopes) are bombarded with neutrons, splitting the uranium nucleus in half and releasing huge amounts of energy. As the nucleus is halved, it emits several extra neutrons, which spin off and split more uranium nuclei, creating still more energy and setting off a chain reaction. This explains the destructive power in an atomic bomb, as well as the constructive power—providing energy to homes and businesses—in a nuclear power plant. Whereas the chain reaction in an atomic bomb becomes an uncontrolled explosion, in a nuclear plant the reaction is slowed and controlled.
Yet nuclear fission is not the most powerful form of atomic reaction. As soon as scientists realized that it was possible to force particles out of a nucleus, they began to wonder if particles could be forced into the nucleus. This type of reaction, known as fusion, puts even nuclear fission, with its awesome capabilities, to shame: nuclear fusion is, after all, the power of the Sun. On the surface of that great star, hydrogen atoms reach incredible temperatures, and their nuclei fuse to create helium. In other words, one element actually transforms into another, releasing enormous amounts of energy in the process.
NUCLEAR ENERGY IN WAR AND PEACE.
The atomic bombs dropped by the United States on Japan in 1945 were fission bombs. These were the creation of a group of scientists—legendary figures such as American physicist J. Robert Oppenheimer (1904-1967), American mathematician John von Neumann (1903-1957), American physicist Edward Teller (1908-), and Italian physicist Enrico Fermi (1901-1954)—involved in the Manhattan Project at Las Alamos, New Mexico.
Some of these geniuses, particularly Oppenheimer, were ambivalent about the moral implications of the enormous destructive power they created. However, most military historians believe that far more lives—both Japanese and American—would have been lost if America had been forced to conduct a land invasion of Japan. As it was, the Japanese surrendered shortly after the cities of Hiroshima and Nagasaki suffered the devastating effects of fission-based explosions.
By 1952, U.S. scientists had developed a "hydrogen," or fusion bomb, thus raising the stakes greatly. This was a bomb that possessed far more destructive capability than the ones dropped over Japan. Fortunately, the Hiroshima and Nagasaki bombs were the only ones dropped in wartime, and a ban on atmospheric nuclear testing has greatly reduced the chances of human exposure to nuclear fallout of any kind. With the end of the arms race between the United States and the Soviet Union, the threat of nuclear destruction has receded somewhat—though it will perhaps always be a part of human life.
Nonetheless, fear of nuclear power, spawned as a result of the arms race, continues to cloud the future of nuclear plants that generate electricity—even though these, in fact, emit less radioactive pollution than coalor gas-burning power plants. At the same time, scientists continue to work on developing a process of power generation by means of nuclear fusion, which, if and when it is achieved, will be one of the great miracles of science.
One of the tools used by scientists researching nuclear fusion is the particle accelerator, which moves streams of charged particles—protons, for instance—faster and faster. These fast particles are then aimed at a thin plate composed of a light element, such as lithium. If the proton manages to be "captured" in the nucleus of a lithium atom, the resulting nucleus is unstable, and breaks into alpha particles.
This method of induced radioactivity is among the most oft-used means of studying nuclear structure and subatomic particles. In 1932, the same year that Chadwick discovered the neutron, English physicist John D. Cockcroft (1897-1967) and Irish physicist Ernest Walton (1903-1995) built the first particle accelerator. Some particle accelerators today race the particles in long straight lines or, to save space, in ringed paths several miles in diameter.
Quantum Theory and Beyond
THE CONTRIBUTION OF RELATIVITY.
It may seem strange that in this lengthy (though, in fact, quite abbreviated!) overview of developments in understanding of the atom, no mention has been made of the figure most associated with the atom in the popular mind: German-American physicist Albert Einstein (1879-1955). The reasons for this are several. Einstein's relativity theory addresses physical, rather than chemical, processes, and did not directly contribute to enhanced understanding of atomic structure or elements. The heart of relativity theory is the famous formula E = mc2, which means that every item of matter possesses energy proportional to its mass multiplied by the squared speed of light.
The value of mc2, of course, is an enormous amount of energy, and in order to be released in significant quantities, an article of matter must experience the kinetic energy associated with very, very high speeds—speeds close to that of light. Obviously, the easiest thing to accelerate to such a speed is an atom, and hence, nuclear energy is a result of Einstein's famous equation. Nonetheless, it should be stressed that although Einstein is associated with unlocking the power of the atom, he did little to explain what atoms are.
However, in the course of developing his relativity theory in 1905, Einstein put to rest a question about atoms and molecules that still remained unsettled after more than a century. Einstein's analysis of Brownian motion, combined with the confirmation of his results by French physicist Jean Baptiste Perrin (1870-1942), showed conclusively that yes, atoms and molecules do exist. It may seem amazing that as recently as 1905, this was still in doubt; however, this only serves to illustrate the arduous path scientists must tread in developing a theory that accurately explains the world.
PLANCK'S QUANTUM THEORY.
A figure whose name deserves to be as much a household word as Einstein's—though it is not—is German physicist Max Planck (1858-1947). It was Planck who initiated the quantum theory that Einstein developed further, a theory that prevails today in the physical sciences.
At the atomic level, Planck showed, energy is emitted in tiny packets or "quanta." Each of these energy packets is indivisible, and the behavior of quanta redefine the old rules of physics handed down from Newton and Maxwell. Thus, it is Planck's quantum theory, rather than Einstein's relativity, that truly marks the watershed, or "before and after," between classical physics and modern physics.
Quantum theory is important not only to physics, but to chemistry as well. It helps to explain the energy levels of electrons, which are not continuous, as in a spectrum, but jump between certain discrete points. The quantum model is now also applied to the overall behavior of the electron; but before this could be fully achieved, scientists had to develop a new understanding of the way electrons move around the nucleus.
BOHR'S PLANETARY MODEL OF THE ATOM.
As was often the case in the history of the atom, a man otherwise respected as a great scientist put forward a theory of atomic structure that at first seemed convincing, but ultimately turned out to be inaccurate. In this case, it was Danish physicist Niels Bohr (1885-1962), a seminal figure in the development of nuclear fission.
Using the observation, derived from quantum theory, that electrons only occupied specific energy levels, Bohr hypothesized that electrons orbited around a nucleus in the same way that planets orbit the Sun. There is no reason to believe that Bohr formed this hypothesis for any sentimental reasons—though, of course, scientists are just as capable of prejudice as anyone. His work was based on his studies; nonetheless, it is easy to see how this model seemed appealing, showing as it did an order at the subatomic level reflecting an order in the heavens.
Many people today who are not scientifically trained continue to think that an atom is structured much like the Solar System. This image is reinforced by symbolism, inherited from the 1950s, that represents "nuclear power" by showing a dot (the nucleus) surrounded by ovals at angles to one another, representing the orbital paths of electrons. However, by the 1950s, this model of the atom had already been overturned.
In 1923, French physicist Louis de Broglie (1892-1987) introduced the particle-wave hypothesis, which indicated that electrons could sometimes have the properties of waves—an eventuality not encompassed in the Bohr model. It became clear that though Bohr was correct in maintaining that electrons occupy specific energy levels, his planetary model was inadequate for explaining the behavior of electrons.
Two years later, in 1925, German physicist Werner Heisenberg (1901-1976) introduced what came to be known as the Heisenberg Uncertainty Principle, showing that the precise position and speed of an electron cannot be known at the same time. Austrian physicist Erwin Schrödinger (1887-1961) developed an equation for calculating how an electron with a certain energy moves, identifying regions in an atom where an electron possessing a certain energy level is likely to be. Schrödinger's equation cannot, however, identify the location exactly.
Rather than being called orbits, which suggest the orderly pattern of Bohr's model, Schrödinger's regions of probability are called orbitals. Moving within these orbitals, electrons describe the shape of a cloud, as discussed much earlier in this essay; as a result, the "electron cloud" theory prevails today. This theory incorporates aspects of Bohr's model, inasmuch as electrons move from one orbital to another by absorbing or emitting a quantum of energy.
WHERE TO LEARN MORE
"The Atom." Thinkquest (Web site). <http://library.thinkquest.org/17940/texts/atom/atom.html> (May 18, 2001).
"Elements" (Web site). <http://home.school.net.hk/~chem/main/F5notes/atom/element.html> (May 18, 2001).
"Explore the Atom" CERN—European Organization for Nuclear Research (Web site). <http://public.web.cern.ch/Public/SCIENCE/Welcome.html> (May 18, 2001).
Gallant, Roy A. The Ever-Changing Atom. New York: Benchmark Books, 1999.
Goldstein, Natalie. The Nature of the Atom. New York: Rosen Publishing Group, 2001.
"A Look Inside the Atom" (Web site). <http://www.aip.org/history/electron/jjhome.htm> (May 18, 2001).
"Portrait of the Atom" (Web site). <http://www.inetarena.com/~pdx4d/snelson/Portrait.html> (May 18, 2001).
"A Science Odyssey: You Try It: Atom Builder." PBS—Public Broadcasting System (Web site). <http://www.pbs.org/wgbh/aso/tryit/atom/> (May 18, 2001).
Spangenburg, Ray and Diane K. Moser. The History of Science in the Nineteenth Century. New York: Facts on File, 1994.
Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.
The smallest particle of an element that retains all the chemical and physical properties of the element. Anatom can exist either alone or in combination with other atoms in a molecule. Atoms are made up of protons, neutrons, and electrons.
ATOMIC MASS UNIT:
An SI unit (abbreviated amu), equal to 1.66 · 10−24 g, for measuring the mass of atoms.
The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table of elements in order of atomic number.
AVERAGE ATOMIC MASS:
A figure used by chemists to specify the mass—in atomic mass units—of the average atom in a large sample.
A one-or two-letter abbreviation for the name of an element.
A substance made up of atoms of more than one element. These atoms are usually joined in molecules.
Negatively charged particles in an atom. Electrons, which spin around the protons and neutrons that make up the atom's nucleus, constitute a very small portion of the atom's mass. The number of electrons and protons is the same, thus canceling out one another; on the other hand, if an atom loses or gainselectrons, it becomes an ion.
Another term for chemical symbol.
An atom or atoms that has lost or gained one or more electrons, and thus has a net electric charge.
Atoms that have an equal number of protons, and hence are of the same element, but differ in their number of neutrons.
The sum of protons and neutrons in an atom's nucleus.
A group of atoms, usually (but not always) representing more than one element, joined in a structure. Compounds are typically made up of molecules.
A subatomic particle that has no electric charge. Neutrons are found at the nucleus of an atom, alongside protons.
A generic term for the heavy particles—protons and neutrons—that make up the nucleus of an atom.
Another term for mass number.
The center of an atom, a region where protons and neutrons are located, and around which electrons spin.
PERIODIC TABLE OF ELEMENTS:
A chart that shows the elements arranged in order of atomic number, along with chemical symbol and the average atomic mass (in atomic mass units) for that particular element. Vertical columns within the periodic table indicate groups or "families" of elements with similar chemical characteristics.
A positively charged particle in an atom. Protons and neutrons, which together form the nucleus around which electrons spin, have approximately the same mass—a mass that is many times greater than that of an electron.
A particle believed to be a component of protons and neutrons. A quarkmay either have a positive electric charge of less than 1+, in which case it is called an "up quark"; or a negative charge of less than 1−, in which case it is called a "down quark."
In a general sense, radiation can refer to anything that travels in astream, whether that stream be composed of subatomic particles or electromagnetic waves. In a more specific sense, the term relates to the radiation from radio active materials, which can be harmful to humanbeings.
A term describing a phenomenon whereby certain isotopes are subject to a form of decay brought about by the emission of high-energy particles or radiation, such as alpha particles, beta particles, or gamma rays.
"Atoms." Science of Everyday Things. . Encyclopedia.com. (December 17, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/atoms-0
"Atoms." Science of Everyday Things. . Retrieved December 17, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/atoms-0
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 478b.c. Our knowledge of the atomic theory of Leucippus is derived almost entirely from the writings of his student, Democritus, who lived around 420b.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 (384–322b.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.
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 (SiO2), alumina (Al2O3), 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 Cu2O. 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.
DEMOCRITUS OF ABBERA
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 H2O. 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 H2O.
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 Hg2O, 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
|DALTON'S FIRST SET OF ATOMIC WEIGHT VALUES (1805)|
|Gaseous oxide of carbon||9.8|
|Carburetted hydrogen from stagnant water||6.3|
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 (NH3) combines exactly with 1 volume of hydrogen chloride gas (HCl) to form a salt (NH4Cl), it is natural to conclude that each volume of gas must contain the same number of particles.
|LAW OF MULTIPLE PROPORTIONS|
|Assumed formula||HgO||Hg 2O|
|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 (O2). 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 (H2). 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 (H2) 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 × 2N is the total
|CANNIZZARO'S METHOD OF MOLECULAR MASS DETERMINATION|
|Gaseous Substance||Density g/L1||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|
|1Density reported for conditions of 100°C and one atmosphere pressure|
|Hydrogen||0.0659||2.00||100||2.00||2||H2||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||H2O||O = 15,88|
|Ammonia||0.557||16.90||17.7||3.00||3||NH3||N = 13.90|
|Methane||0.524||15.90||25.1||4.00||4||CH4||C = 11.90|
mass of the hydrogen atoms in the 1 liter sample of hydrogen molecules; recall that hydrogen was shown to be diatomic (H2) 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 HClN, 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
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.
"Atoms." Chemistry: Foundations and Applications. . Encyclopedia.com. (December 17, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/atoms
"Atoms." Chemistry: Foundations and Applications. . Retrieved December 17, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/atoms
atom [Gr.,=uncuttable (indivisible)], basic unit of matter; more properly, the smallest unit of a chemical element having the properties of that element.
Structure of the Atom
The atom consists of a central, positively charged core, the nucleus, and negatively charged particles called electrons that are found in orbits around the nucleus.
Almost the entire mass of the atom is concentrated in the nucleus, which occupies only a tiny fraction of the atom's volume. The nucleus of an atom consists of neutrons and protons, the neutron being an uncharged particle and the proton a positively charged one. Their masses are almost equal. Atoms containing the same number of protons but different numbers of neutrons represent different forms, or isotopes, of the same element.
Surrounding the nucleus of an atom are its electrons; for a neutral atom, the number of electrons is equal to the atomic number. The outermost electrons of an atom determine its chemical and electrical properties. An atom may combine chemically with another atom in various ways, either by giving up or receiving electrons, thus setting up an electrical attraction between the atoms (see ion), or by sharing one or more pairs of electrons (see chemical bond). Because metals have few outermost electrons and tend to give them up easily, they are good conductors of electricity or heat (see conduction).
The electrons are often described as revolving about the nucleus as the planets revolve about the sun. This picture, however, is misleading. The quantum theory has shown that all particles in motion also have certain wave properties. For a particle the size of an electron, these properties are of considerable importance. As a result the electrons in an atom cannot be pictured as localized in space, but rather should be viewed as smeared out over the entire orbit so that they form a cloud of charge. The electron clouds around the nucleus represent regions in which the electrons are most likely to be found. The shapes of these clouds can be very complex, in marked contrast to the simple elliptical orbits of planets. Surprisingly, the sizes of all atoms are comparable, in spite of the large differences in the number of electrons they contain.
Atomic Weight and Number
The atomic number of an atom is simply the number of protons in its nucleus. The atomic weight of an atom is given in most cases by the mass number of the atom, equal to the total number of protons and neutrons combined. An atom may be conveniently symbolized by its chemical symbol with the atomic number and mass number written as subscript and superscript, respectively. For example, the symbol for uranium is U (atomic number 92); the isotopes of uranium with atomic weights 235 and 238 are indicated by 23592U and 23892U.
Development of Atomic Theory
Early Atomic Theory
The atomic theory, which holds that matter is composed of tiny, indivisible particles in constant motion, was proposed in the 5th cent. BC by the Greek philosophers Leucippus and Democritus and was adopted by the Roman Lucretius. However, Aristotle did not accept the theory, and it was ignored for many centuries. Interest in the atomic theory was revived during the 18th cent. following work on the nature and behavior of gases (see gas laws).
From Dalton to the Periodic Table
Modern atomic theory begins with the work of John Dalton, published in 1808. He held that all the atoms of an element are of exactly the same size and weight (see atomic weight) and are in these two respects unlike the atoms of any other element. He stated that atoms of the elements unite chemically in simple numerical ratios to form compounds. The best evidence for his theory was the experimentally verified law of simple multiple proportions, which gives a relation between the weights of two elements that combine to form different compounds.
Evidence for Dalton's theory also came from Michael Faraday's law of electrolysis. A major development was the periodic table, devised simultaneously by Dmitri Mendeleev and J. L. Meyer, which arranged atoms of different elements in order of increasing atomic weight so that elements with similar chemical properties fell into groups. By the end of the 19th cent. it was generally accepted that matter is composed of atoms that combine to form molecules.
Discovery of the Atom's Structure
In 1911, Ernest Rutherford developed the first coherent explanation of the structure of an atom. Using alpha particles emitted by radioactive atoms, he showed that the atom consists of a central, positively charged core, the nucleus, and negatively charged particles called electrons that orbit the nucleus. There was one serious obstacle to acceptance of the nuclear atom, however. According to classical theory, as the electrons orbit about the nucleus, they are continuously being accelerated (see acceleration), and all accelerated charges radiate electromagnetic energy. Thus, they should lose their energy and spiral into the nucleus.
This difficulty was solved by Niels Bohr (1913), who applied the quantum theory developed by Max Planck and Albert Einstein to the problem of atomic structure. Bohr proposed that electrons could circle a nucleus without radiating energy only in orbits for which their orbital angular momentum was an integral multiple of Planck's constant h divided by 2π. The discrete spectral lines (see spectrum) emitted by each element were produced by electrons dropping from allowed orbits of higher energy to those of lower energy, the frequency of the photon of light emitted being proportional to the energy difference between the orbits.
Around the same time, experiments on x-ray spectra (see X ray) by H. G. J. Moseley showed that each nucleus was characterized by an atomic number, equal to the number of unit positive charges associated with it. By rearranging the periodic table according to atomic number rather than atomic weight, a more systematic arrangement was obtained. The development of quantum mechanics during the 1920s resulted in a satisfactory explanation for all phenomena related to the role of electrons in atoms and all aspects of their associated spectra. With the discovery of the neutron in 1932 the modern picture of the atom was complete.
Contemporary Studies of the Atom
With many of the problems of individual atomic structure and behavior now solved, attention has turned to both smaller and larger scales. On a smaller scale the atomic nucleus is being studied in order to determine the details of its structure and to develop sources of energy from nuclear fission and fusion (see nuclear energy), for the atom is not at all indivisible, as the ancient philosophers thought, but can undergo a number of possible changes. On a larger scale new discoveries about the behavior of large groups of atoms have been made (see solid-state physics). The question of the basic nature of matter has been carried beyond the atom and now centers on the nature of and relations between the hundreds of elementary particles that have been discovered in addition to the proton, neutron, and electron. Some of these particles have been used to make new types of exotic "atoms" such as positronium (see antiparticle) and muonium (see muon).
See G. Gamow, The Atom and Its Nucleus (1961); H. A. Boorse and L. Motz, ed., The World of the Atom (2 vol., 1966); B. H. Bransden and C. J. Joachain, Physics of Atoms and Molecules (1986).
"atom." The Columbia Encyclopedia, 6th ed.. . Encyclopedia.com. (December 17, 2017). http://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/atom
"atom." The Columbia Encyclopedia, 6th ed.. . Retrieved December 17, 2017 from Encyclopedia.com: http://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/atom
Atoms and the subatomic particles that comprise them, are the elementary building blocks of material substances. Although the term atom, derived from the Greek word atomos, meaning indivisible, would seem inappropriate for an entity that, as science has established, is divisible, the word atom still makes sense, because, depending on the context, atoms can still be regarded as indivisible. Namely, once the nucleus is split, the atom loses its identity and subatomic particles. Protons, electrons, neutrons, are all the same—regardless of the type of atom or element—it is only their numbers and unique combinations that make for different atoms. Accordingly, an atom is the smallest particle of an element.
Atoms share many characteristics with other material objects: they can be measured, and they also have mass and weight. Because traditional methods of measuring are difficult to use for atoms and subatomic particles, scientists have created a new unit, the atomic mass unit (amu), which is defined as the one-twelfth of the mass of the average carbon atom.
The principal subatomic particles are the protons, neutrons, and the electrons. The nucleus, the atom's core, consists of protons, which are positively charged particles, and neutrons, particles without any charge. Electrons are negatively charged particles with negligible mass that orbit around the nucleus. An electron's mass is so small that it is usually given a 0 amu value in atomic mass units, compared to the value of 1 amu assigned to neutrons and protons (neutrons do carry slightly more mass than protons and neither exactly equals 1 amu—but for purposes of this article the approximate values will suffice). In fact, as the nucleus represents more than 99% of an atom's mass, it is interesting to note that an atom is mostly space . For example, if a hydrogen atom's nucleus were enlarged to the size of a marble , the atom's diameter (to the electron orbit) would be around 0.5 mi (800 m).
At one time, scientists asserted that electrons circled around the nucleus in planet-like orbits. However, because all subatomic particles, including electrons, exhibit wave-like properties, it is makes no sense to conceptualize the movement of electrons as like planetary rotation . Scientists therefore prefer terms like "electron cloud patterns," or "shells," indicating an electron's position and/or pattern of movement in relation to the nucleus. Thus, for example, hydrogen has one electron in its innermost, lowest energy shell (a shell is also an energy level); lithium—with three electrons—has two shells, with inner most, lowest energy shell contains two electrons that one electron exists in a more distant shell or higher potential energy level. The elements exhibit four distinctive shapes of shell—designated s, p, d, and f orbitals.
While subatomic particles are generic and interchangeable, in combination they determine an atom's identity. For example, we know that an atom with a nucleus consisting of one proton must be hydrogen (H). An atom with two protons is always a helium (He) atom. Thus, we see that the key to an atom's identity is to be found in the atom's inner structure. In addition, a electrically balanced chemical element is an instance of atomic electronic equilibrium: for example, in an electrically balanced chemical element, the number of positively charged particles (protons) always equals the number of negatively charged particles (electrons). A loss or gain of electrons results in a net charge and the atom becomes an ion.
Although the number of protons determines the name (type) of atom, each atom may be heavier or lighter depending on the number of neutrons present. Atoms of the same element with different mass (reflecting differing numbers of neutrons) are isotopes.
Research into the atom's nucleus has uncovered a variety of subatomic particles, including quarks and gluons. Considered by some researchers the true building blocks of matter, quarks are the particles that form protons and neutrons. Gluons hold smaller clusters of quarks together.
The atom is best characterized by the laws and terminology or quantum physics . On a larger scale, chemists study reactions, the behavior of elements in interaction, and reactions, such as those leading to the formation of chemical compounds. Such reactions involve the transfer of electrons and/or the sharing of electrons in atomic bonds.
For example, the formation of sodium chloride, also known as table salt, would be impossible without specific changes at a subatomic level. The genesis of sodium chloride (NaCl) starts when a sodium (Na) atom, which has 11 electrons, loses an electron. With 10 electrons, the atom now has one more proton than electrons and thus becomes a net positively charged sodium ion Na+ (a positively charged ion is also known as a cation. Chlorine becomes a negatively charged anion by accepting a free electron to take on a net negative charge. The newly acquired electron goes into the outer shell, also known as the valence shell that already contains seven electrons. The addition of the eighth electron to the chlorine atom's outmost shell fulfills the octet rule and allows the atom—although now a negatively charged chlorine ion (Cl−)—to be more stable. The electrical attraction of the sodium cations for the chlorine anion results in an ionic bond to form salt. Crystals of table salt consist of equal numbers of sodium cations and chlorine anions, cation-anion pairs being held together by a force of electrical attraction.
The octet rule is used to describe the attraction of elements toward having, whenever possible, eight valence-shell electrons (four electron pairs) in their outer shell. Because a full outer shell with eight electrons is relatively stable, many atoms lose or gain electrons to obtain an electron configuration like that of the nearest noble gas. Except for helium (with a filled 1s shell), noble gases have eight electrons in their valence shells.
Interestingly, not long after scientists realized that at the level of the nucleus an atom is divisible, transmutation, or the old alchemic dream of turning one substance into another, became a reality. Fission and fusion are tranformative processes that, by altering the nucleus, alter the element. For example, scientists even succeeded in creating gold by bombarding platinum-198 with neutrons to create platinum-199 that then decays to gold-199. Although clearly demonstrating the reality of transmutation, this particular transmutation (a change in the nuclear structure that changes one element into another) is by no means an easy or cheap method of producing gold. Quite the contrary, because platinum, particularly the platinum-199 isotope, is more expensive than gold produced. Regardless, the symbolic value of the experiment is immense, as it shows that the idea, developed by ancient alchemists and philosophers, of material transmutation—accomplished at the nuclear level—does not essentially contradict our understanding of the atom.
Natural transformations also exist—as with the decay of Carbon-14 to nitrogen—accomplished by the nuclear transformation of one Carbon-14 neutron into a proton.
See also Atomic mass and weight; Atomic number; Atomic theory; Chemical bonds and physical properties; Chemical elements
"Atom." World of Earth Science. . Encyclopedia.com. (December 17, 2017). http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/atom
"Atom." World of Earth Science. . Retrieved December 17, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/atom
An atom is the smallest particle of a element that has all the properties of that element. Imagine that you decide to cut a chunk of aluminum metal into half, over and over again. At some point, you would need very small tools to do the cutting, tools smaller than anything that really exists. However, you would eventually get to the very smallest piece of aluminum that still has all the properties of the original chunk. That smallest piece is an atom of aluminum.
One of the questions that ancient Greeks thinkers debated was the structure of matter. Is matter, they asked, continuous or discontinuous? That is, in the aluminum example mentioned above, can a person continue to cut a chunk of aluminum into smaller pieces for ever and ever? Or would the person eventually reach some smallest piece of aluminum that could be divided no further?
Two of the philosophers who argued for the latter opinion were Leucippus (born about 490 b.c.) and his student Democritus (c. 470–c. 380 b.c.). It was Democritus, in fact, who first used the word atomos to describe the smallest possible particles of matter. Atomos means "indivisible" in Greek.
The particle theory of matter was not developed to any great extent for more than 2,000 years. Then, in 1808, English chemist John Dalton (1766–1844) rephrased the theory in modern terms. Dalton thought of atoms as tiny, indivisible particles, similar to ball bearings or marbles. Dalton's theory of atoms satisfactorily explained what was then known about matter; it was quickly accepted by many other (although not all) chemists.
In the two centuries since Dalton first proposed the modern concept of atoms, that concept has undergone some dramatic changes. We no longer believe that atoms are indivisible particles. We know that they consist of smaller units, known as protons, neutrons, and electrons. These particles are called subatomic particles because they are all smaller than an atom itself. Some subatomic particles are capable of being divided into even smaller units known as quarks.
Modern models of the atom
Scientists think of atoms today in mathematical terms. They use mathematical equations to represent the likelihood of finding electrons in various parts of the atom and to describe the structure of the atomic nucleus, in which protons and neutrons exist.
Most people still find it helpful to think about atoms in physical terms that we can picture in our minds. For most purposes, these pictures are good enough to understand what atoms are like. An atom consists of two parts, a nucleus and a set of one or more electrons spinning around the nucleus.
The nucleus is located at the center of an atom. It consists of one or more protons and, with the exception of the hydrogen atom, one or more neutrons. The number of protons in an atom is given the name atomic number. An atom with one proton in its nucleus has an atomic number of 1, while an atom with sixteen protons in its nucleus has an atomic number of 16. The total number of protons and neutrons in a nucleus is called the atom's mass number. An atom with two protons and two neutrons, for example, has a mass number of 4.
The number of electrons located outside the nucleus of an atom is always the same as the number of protons. An atom with seven protons in its nucleus (no matter how many neutrons) also has seven electrons outside the nucleus. Those electrons travel in paths around the nucleus somewhat similar to the orbits followed by planets around the Sun. Each
of these orbits can hold a certain number of electrons. The first orbit, for example, may hold up to two electrons, but no more. The second orbit may hold up to eight electrons, but no more. The third orbit may hold a maximum of 18 electrons.
These limits determine how the electrons in an atom are distributed. Suppose that the nucleus of an atom contains nine protons. Then the atom also contains nine electrons outside the nucleus. Two of the electrons can be in the first orbit around the nucleus, but the other seven must go to the second orbit.
The term electron orbit is not really correct, even if it does help understand what an electron's path looks like. A better term is electron energy level. The closer an electron is to the nucleus of an atom, the less energy it has; the farther away from the nucleus, the more energy it has.
An atom and the particles of which it is composed can be fully described by knowing three properties: mass, electrical charge, and spin.
The mass of protons, neutrons, and electrons is so small that normal units of measurement (such as the gram or centigram) are not used. As an example, the actual mass of a proton is 1.6753 × 10−24 g, or 0.000 000 000 000 000 000 000 001 675 3 grams. Numbers of this size are so inconvenient to work with that scientists have invented a special unit known as the atomic mass unit (abbreviation: amu) to state the mass of subatomic particles. One atomic mass unit (1 amu) is approximately equal to the mass of a single proton. Using this measure, the mass of a neutron is also about 1 amu, and the mass of an electron, about 0.00055 amu.
The mass of an atom, then, is equal to the total mass of all protons, neutrons, and electrons added together. In the case of the oxygen atom, that mass is expressed as follows:
mass of oxygen atom = mass of 8 protons + 8 neutrons + 8 electrons
mass of one oxygen atom = 8 amu + 8 amu + (8 × 0.00055 amu)
mass of one oxygen atom = 16.0044 amu
The total mass of an atom is called its atomic mass or, less accurately, its atomic weight. As you can see, the mass of an atom depends primarily on the mass of its protons and neutrons and is hardly affected by the mass of its electrons.
The actual mass and size of atoms using ordinary units of measurement are both very small. The mass of one oxygen atom measured in grams is 5.36 × 10−23 g or 0.000 000 000 000 000 000 000 053 6 grams.
The dimensions of an atom and its nucleus are also amazingly small. The distance across the outside of a typical atom is about 10−10 m, or 0.000 000 000 1 meters. In contrast, the distance across a nucleus is about 10−15 m, or 0.000 000 000 000 001 meters. In another words, an atom is about 100,000 times larger in size than it its nucleus. To get some idea of this comparison, imagine a pea placed in the center of a large football stadium. If the pea represents the nucleus of an atom, the closest electrons in the atom would be spinning around outside the outermost reaches of the stadium's upper seats.
[See also Atomic mass; Atomic theory; Electron; Element, chemical; Matter, states of; Subatomic particles ]
"Atom." UXL Encyclopedia of Science. . Encyclopedia.com. (December 17, 2017). http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/atom-1
"Atom." UXL Encyclopedia of Science. . Retrieved December 17, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/atom-1
"atom." World Encyclopedia. . Encyclopedia.com. (December 17, 2017). http://www.encyclopedia.com/environment/encyclopedias-almanacs-transcripts-and-maps/atom
"atom." World Encyclopedia. . Retrieved December 17, 2017 from Encyclopedia.com: http://www.encyclopedia.com/environment/encyclopedias-almanacs-transcripts-and-maps/atom