It is impossible to see the history of the periodic table and the discovery of the elements separate from the history of the atom and its particles and subparticles. Therefore this section will try to provide a short introduction into its long history as well as a description of its basic properties.
1.2.1 History
The concept that matter consists of discrete units is a very old idea, appearing in many ancient cultures such as Greece and India. The word atomos, meaning “uncuttable,” was “invented” by the ancient Greek philosophers Leucippus and his student Democritus (c.460 370 BCE) (Fig. 1.10). Democritus taught that atoms were infinite in number, uncreated, and eternal, and that the qualities of an object result from the kind of atoms that compose it. Democritus’s atomism was refined and expanded by the later philosopher Epicurus (341 270 BCE). In the period of the Early Middle Ages, atom- ism was mostly forgotten in western Europe, but survived among some groups of Islamic philosophers. During the 12th century, atomism became known again in western Europe because of references found in the newly rediscovered writ- ings of Aristotle. In the 14th century, the rediscovery of major works describing atomist teachings, including the Roman poet and philosopher Titus Lucretius Carus’ De rerum natura (c. October 15, 99 BCE to c.55 BCE) and biogra- pher of Greek philosophers Diogenes Lae¨rtius’s Lives and Opinions of Eminent Philosophers, caused an increased scholarly attention on the subject. However, since atomism was linked to the philosophy of Epicureanism, which con- tradicted orthodox Christian teachings, belief in atoms was not considered acceptable. The French Catholic priest, phi- losopher, astronomer, and mathematician Pierre Gassendi (January 22, 1592 to October 24, 1655) revived Epicurean atomism with amendments, stating that atoms were created by God and, though extremely numerous, are not infinite.
His modified theory of atoms was spread in France by the physician and traveler Francáois Bernier (September 25, 1620 to September 22, 1688) and in England by the natural philosopher and writer Walter Charleton (February 2, 1619 to April 24, 1707). The Anglo-Irish natural philosopher, chemist, physicist, and inventor Robert Boyle (January 25, 1627 to December 31, 1691) and the English mathematician, physicist, astronomer, theologian, and author Isaac Newton (December 25, 1642 to March 20, 1727) both defended atomism and, by the end of the 17th century, it was accepted by parts of the scientific community.
In the early 1800s English chemist, physicist, and meteorologist John Dalton (September 6, 1766 to July 27, 1844) (Fig. 1.11) used the idea of atoms to describe why elements always react in ratios of small whole numbers (the law of multiple proportions). For instance, there are two types of tin oxide: one is 88.1% tin and 11.9% oxygen and the other is 78.7% tin and 21.3% oxygen [tin(II) oxide and tin dioxide, respectively]. This means that 100 g of tin can either
FIGURE 1.10 Democritus.
combine with 13.5 or 27 g of oxygen. 13.5 and 27 form a ratio of 1:2, a ratio of small whole numbers. This general pat- tern in chemistry indicated to Dalton that elements react in multiples of discrete units—in other words, atoms. In the case of tin oxides, one tin atom will combine with either one or two oxygen atoms. Dalton also held the belief that the atomic theory could explain why water absorbs different gases in different proportions. For example, he found that water absorbs carbon dioxide far better than it absorbs nitrogen. Dalton hypothesized that this was due to the differences between the masses and configurations of the gases’ respective particles, and carbon dioxide molecules (CO2) are heavier and larger than nitrogen molecules (N2).
In 1827, Scottish botanist and paleobotanist Robert Brown (December 21, 1773 to June 10, 1858) used a microscope to observe dust grains floating in water and discovered that they moved about erratically, a phenomenon that is now known as “Brownian motion.” This was thought to be caused by water molecules knocking the grains about. In 1905 German-born theoretical physicist Albert Einstein (March 14, 1879 to April 18, 1955) demonstrated the reality of these molecules and their motions by producing the first statistical physics analysis of Brownian motion. French physicist Jean Perrin (September 30, 1870 to April 17, 1942) applied Einstein’s work to experimentally measure the mass and dimensions of atoms, thereby definitively confirming Dalton’s atomic theory.
The physicist Sir Joseph John Thomson (December 18, 1856 to August 30, 1940) determined the mass of cathode rays, showing they consisted of particles, but were about 1800 times lighter than the lightest atom, hydrogen(Thomson, 1901) (Fig. 1.12). Hence, they could not be atoms, but a new particle, the first subatomic particle to be discovered, which he initially called “corpuscle” but was later renamed electron, after particles postulated by Irish physicist George Johnstone Stoney (February 15, 1826 to July 5, 1911) in 1874. Thomson also proved that they were identical to parti- cles given off by photoelectric and radioactive materials. It was soon recognized that they are the particles that carry electric currents in metal wires and carry the negative electric charge within atoms. He received the 1906 Nobel Prize in Physics for this work. With this work he overturned the belief that atoms are the indivisible, ultimate particles of matter. In addition, Thomson erroneously postulated that the low mass, negatively charged electrons were distributed throughout the atom in a uniform sea of positive charge. This became known as the plum pudding model.
In 1909 German physicist Johannes Wilhelm “Hans” Geiger (September 30, 1882 to September 24, 1945) and English-New Zealand physicist Ernest Marsden (February 19, 1889 to December 15, 1970), under the direction of New Zealand physicist Ernest Rutherford (August 30, 1871 to October 19, 1937), bombarded a metal foil with alpha
FIGURE 1.11 John Dalton by Charles Turner (1773 1857) after James Lonsdale (1777 1839) (Mezzotint).
particles (which we now know consists of two protons and two neutrons bound together into a particle identical to a helium-4 nucleus) to determine how they scattered. They anticipated all the alpha particles to pass straight through with little deflection, since Thomson’s model predicted that the charges in the atom are so diffuse that their electric fields should not affect the alpha particles much. Nevertheless, Geiger and Marsden observed alpha particles being deflected by angles greater than 90 degrees, something that was thought to be impossible based on Thomson’s model. To explain this, Rutherford proposed that the positive charge of the atom is concentrated in a tiny nucleus at the center of the atom.
While investigating the products of radioactive decay, in 1913 English radiochemist Frederick Soddy (September 2, 1877 to September 22, 1956) found that there seemed to be more than one type of atom at each position on the periodic table(Soddy, 1913a-c)(Fig. 1.13). He also explained, with Ernest Rutherford, that radioactivity is due to the transmuta- tion of elements, now known to involve nuclear reactions. The term isotope, Greek for at the same place, was suggested by Scottish doctor and writer Margaret Todd (April 23, 1859 to September 3, 1918) as a suitable name for different atoms that belong to the same element. Todd was a family friend of chemist Frederick Soddy, and a lecturer at the University of Glasgow. In 1913 Soddy explained to her the research on radioactivity for which he later won the Nobel Prize in Chemistry in 1921. The term isotope was accepted and used by Soddy and has become standard scientific nomenclature. J.J. Thomson created a technique for isotope separation through his work on ionized gases, which subse- quently led to the discovery of stable isotopes [The term stable isotope has a meaning similar to stable nuclide but is preferably used when speaking of nuclides of a specific element. Therefore the plural form “stable isotopes” usually refers to isotopes of the same element. The relative abundance of such stable isotopes can be determined experimentally (isotope analysis), resulting in an isotope ratio that can be used as a research tool. Theoretically, such stable isotopes can include the radiogenic daughter products of radioactive decay, used in radiometric dating. Nevertheless, the expres- sion stable-isotope ratio is preferably used to refer to isotopes whose relative abundances are affected by isotope frac- tionation in nature. This field is known as stable-isotope geochemistry.].
In 1913 the Danish physicist Niels Bohr (October 7, 1885 to November 18, 1962) (Fig. 1.14) developed a model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits (similar to planets around the sun) and could jump between these orbits only with discrete changes of energy corresponding to absorption or radiation of a photon (Fig. 1.15). This quantization was used to explain why the electron orbits are
FIGURE 1.12 J.J. Thomson (pre-1915).
stable (given that normally, charges in acceleration, including circular motion, lose kinetic energy that is emitted as electromagnetic radiation) and why elements absorb and emit electromagnetic radiation in discrete spectra (Bohr, 1922a,b). Later in that same year English physicist Henry Gwyn Jeffreys Moseley (November 23, 1887 to August 10, 1915) showed additional experimental proof supporting Niels Bohr’s theory. These results refined Ernest Rutherford’s and Dutch lawyer and amateur physicist Antonius Van den Broek’s model (May 4, 1870 to October 25, 1926), which suggested that the atom contains in its nucleus a number of positive nuclear charges that is equal to its (atomic) number
FIGURE 1.13 Frederick Soddy [Nobel Prize in Chemistry (1921)].
FIGURE 1.14 Niels Bohr (c.1922).
in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quan- tity. That it is equal to the atomic nuclear charge is still the accepted atomic model today.
Chemical bonds between atoms could now be explained, by American physical chemist and a former Dean of the College of Chemistry at University of California, Berkeley Gilbert Newton Lewis [October 25 (or 23), 1875 to March 23, 1946] in 1916, as the interactions between their constituent electrons(Lewis, 1916). Since it was known that the chemical properties of the elements largely repeat themselves consistent with the periodic law, in 1919 the American chemist and physicist Irving Langmuir (January 31, 1881 to August 16, 1957) proposed that this could be explained if the electrons in an atom were connected or clustered in some manner (Langmuir, 1919a,b). Groups of electrons were believed to occupy a set of electron shells around the nucleus [In chemistry and atomic physics, an electron shell, or a principal energy level, may be thought of as an orbit followed by electrons around an atom’s nucleus. The closest shell to the nucleus is called the “1 shell” (also called “K shell”), followed by the “2 shell” (or “L shell”), then the “3 shell”
(or “M shell”), and so on farther and farther from the nucleus. The shells correspond with the principal quantum num- bers (n51, 2, 3, 4. . .) or are labeled alphabetically with letters used in the X-ray notation (K, L, M,. . .). Each shell can contain only a fixed number of electrons: the first shell can hold up to two electrons; the second shell can hold up to eight (216) electrons; the third shell can hold up to 18 (216110); and so on. The general formula is that the nth shell can in principle hold up to 2(n2) electrons. Since electrons are electrically attracted to the nucleus, an atom’s elec- trons will generally occupy outer shells only if the more inner shells have already been completely filled by other elec- trons. However, this is not a strict requirement: atoms may have two or even three incomplete outer shells. The electrons in the outermost occupied shell (or shells) determine the chemical properties of the atom; it is called the valence shell. Each shell consists of one or more subshells, and each subshell consists of one or more atomic orbitals.].
The Stern Gerlach experiment of 1922 conducted by the German-American physicist Otto Stern (February 17, 1888 to August 17, 1969) and German physicist Walter Gerlach (August 1, 1889 to August 10, 1979) provided further proof of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom’s angular momentum, or spin. As this spin direction is initially random, the beam would be expected to deflect in a random direction. Instead, the beam was split into two directional components, corresponding to the atomic spin being oriented up or down with respect to the magnetic field. Historically, this experiment was decisive in convincing physicists of the reality of angular- momentum quantization in all atomic-scale systems. In 1925 German theoretical physicist Werner Heisenberg (December 5, 1901 to February 1, 1976) (Fig. 1.16) reported the first consistent mathematical formulation of quantum mechanics (Matrix Mechanics) (Heisenberg, 1925). One year earlier, in 1924, French physicist Louis Victor Pierre Raymond de Broglie (August 15, 1892 to March 19, 1987) had suggested that all particles behave to an extent like waves and, in 1926, Austrian physicist Erwin Rudolf Josef Alexander Schro¨dinger (August 12, 1887 to January 4, 1961) (Fig. 1.17) used this idea to come up with a mathematical model of the atom (Wave Mechanics) that described the electrons as three-dimensional waveforms rather than point particles (Schro¨dinger, 1926). A consequence of
FIGURE 1.15 Bohr’s model of an atom.
utilizing waveforms to define particles is that it is mathematically impossible to obtain precise values for both the posi- tion and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. This model could explain observations of atomic behavior that previous models were not able to, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus the planetary model of the atom (Bohr’s model) was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed. The advances in mass spectrometry allowed the mass of atoms to be determined with increased accuracy. A mass spectrometer uses a magnet to bend the trajectory
FIGURE 1.16 Werner Heisenberg.Bundesarchiv, Bild 183-R57262/Unknown/CC-BY-SA 3.0.
FIGURE 1.17 Erwin Schrodinger (1933).
of a beam of ions, and the amount of deflection is determined by the ratio of an atom’s mass to its charge. The English chemist and physicist Francis William Aston (September 1, 1877 to November 20, 1945) proved with this instrument that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole num- ber rule. The explanation for these different isotopes had to wait for the discovery of the neutron, an uncharged particle with a mass similar to the proton, by the British physicist James Chadwick (October 20, 1891 to July 24, 1974) in 1932 (Chadwick, 1935)(Fig. 1.18). Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus. In 1938, the German chemist Otto Hahn (March 8, 1879 to July 28, 1968), a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. In its place, his chemical experiments showed barium as a product. A year later, Austrian-Swedish physicist Lise Meitner (November 7, 1878 to October 27, 1968) and her nephew Austrian physicist Otto Frisch (October 1, 1904 to September 22, 1979) con- firmed that Hahn’s results were the first experimental nuclear fission (Meitner and Frisch, 1939). In 1944, Hahn received the Nobel Prize in Chemistry. Despite Hahn’s efforts, the contributions of Meitner and Frisch were not recog- nized. Based on their original correspondence, many historians have documented their view of the discovery of nuclear fission and belief that Meitner should have been awarded the Nobel Prize with Hahn, a classic example of the Matilda effect [a bias against acknowledging the achievements of those women scientists whose work is attributed to their male colleagues. This effect was first described by suffragist and abolitionist Matilda Joslyn Gage (1826 98) in her essay,
“Woman as Inventor.” The term “Matilda effect” was coined in 1993 by science historian Margaret W. Rossiter.]. At the end of World War II in 1945, Hahn was suspected of working on the German nuclear weapon project to develop an atomic reactor or an atomic bomb, but his only connection was the discovery of fission; he did not work on the pro- gram. In April 1945, Hahn and nine leading German physicists (including Max von Laue, Werner Heisenberg, and Carl Friedrich von Weizsa¨cker) were taken into custody by the Alsos Mission and interned at Farm Hall, Godmanchester, near Cambridge, England, from 3 July 1945 to 3 January 1946. Hahn was still being detained at Farm Hall when the announcement was made; thus his whereabouts were a secret, and it was impossible for the Nobel committee to send him a congratulatory telegram. Instead, he learned about his award through the Daily Telegraph newspaper. His fellow interned German scientists celebrated his award on November 18 by giving speeches, making jokes, and composing songs. On December 4, Hahn was persuaded by two of his captors to write a letter to the Nobel committee accepting the prize but also stating that he would not be able to attend the award ceremony. He could not participate in the Nobel festivities on December 10 since his captors would not allow him to leave Farm Hall. In the 1990s, the records of the Nobel Prize committee that decided on that prize were opened. Based on this information, several scientists and journal- ists have called her exclusion “unjust,” and Meitner has received many posthumous honors, including naming chemical element 109 meitnerium in 1992. Despite not having been awarded the Nobel Prize, Lise Meitner was invited to attend
FIGURE 1.18 James Chadwick (c.1945).rCopyright Triad National Security, LLC. All Rights Reserved.
the Lindau Nobel Laureate Meeting in 1962 (annual, scientific conferences held in Lindau, Bavaria, Germany since 1951. Their aim is to bring together Nobel laureates and young scientists to foster scientific exchange between different generations and cultures.). In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks. The standard model of particle physics developed so far has successfully explained the properties of the nucleus in terms of these subatomic particles and the forces that govern their interactions.
1.2.2 Structure and properties
Though the word atom originally represented a particle that cannot be divided into smaller particles, in modern scien- tific usage the atom is composed of a variety of subatomic particles. These fundamental particles of an atom are the electron (e or e2), the proton (p or p1) and the neutron (n); all three are fermions (a particle that follows Fermi Dirac statistics. These particles obey the Pauli exclusion principle. Fermions include all quarks and leptons, as well as all composite particles made of an odd number of these, such as all baryons and many atoms and nuclei. Fermions differ from bosons, which obey Bose Einstein statistics.). Nevertheless, the hydrogen-1 atom has no neutrons and the hydron ion has no electrons [general name for a cationic form of atomic hydrogen, represented with the symbol H1. However, this term is avoided and instead “proton” is used, which strictly speaking refers to the cation of protium, the most com- mon isotope of hydrogen. The term “hydron” includes cations of hydrogen regardless of their isotopic composition;
thus it refers collectively to protons (1H1) for the protium isotope, deuterons (2H1 or D1) for the deuterium isotope, and tritons (3H1 or T1) for the tritium isotope. Unlike most other ions, the hydron consists only of a bare atomic nucleus.]. The electron is the least massive of these particles at 9.11310231kg, with a negative electrical charge and a size that is too small to be measured using currently available techniques. It was the lightest particle with a positive rest mass determined, until the discovery of the neutrino mass. Under normal conditions, electrons are bound to the posi- tively charged nucleus due to the attraction created by opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes, respectively, negatively or positively charged as a whole; a charged atom is called an ion (positive charge2cation, negative charge2anion). Electrons have been known since the late 19th cen- tury, mostly because of J.J. Thomson’s research. Protons have a positive charge and a mass 1836 times that of the elec- tron, at 1.6726310227kg. The number of protons in an atom is known as its atomic number. Ernest Rutherford detected that nitrogen under alpha-particle bombardment ejects what seemed to be hydrogen nuclei. By 1920 he had recognized that the hydrogen nucleus forms a distinct particle within the atom and called it a proton. Neutrons have no electrical charge and have a free mass of 1839 times the mass of the electron, or 1.6749310227kg. Neutrons are the heaviest of the three fundamental particles, but their mass can be reduced by the nuclear binding energy (the minimum energy that would be required to disassemble the nucleus of an atom into its component parts. The binding is always a positive number, as it is necessary to spend energy in moving these nucleons, attracted to each other by the strong nuclear force, away from each other. The mass of an atomic nucleus is less than the sum of the individual masses of the free constituent protons and neutrons, according to Einstein’s equationE5mc2, wheremis the mass loss andcis the speed of light (in vacuum). This “missing mass” is known as the mass defect and characterizes the energy that was released when the nucleus was formed.). Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of 2.5310215m—although the “surface” of these particles is not sharply defined.
In the standard model of physics, electrons are truly elementary particles with no internal structure. In contrast, both protons and neutrons are composite particles consisting of elementary particles called quarks. Quarks combine to form composite particles called hadrons, the most stable of which are protons and neutrons, the components of atomic nuclei.
Due to a phenomenon known as color confinement, quarks are never directly observed or found in isolation; they can be found only within hadrons, which include baryons (such as protons and neutrons) and mesons [In particle physics, mesons are hadronic subatomic particles composed of one quark and one antiquark, bound together by strong interac- tions. Because mesons are composed of quark subparticles, they have physical size, notably a diameter of roughly 1 fm, which is about 1.2 times the size of a proton or neutron. All mesons are unstable, with the longest-lived lasting for only a few hundredths of a microsecond. Charged mesons decay (sometimes through mediating particles) to form electrons and neutrinos. Uncharged mesons may decay to photons. Both of these decays indicate that color is no longer a property of the byproducts.]. Therefore, much of what is known about quarks has been based upon observations of hadrons.
Quarks have several intrinsic properties, including electric charge, mass, color charge, and spin. They are the only ele- mentary particles in the standard model of particle physics to experience all four fundamental interactions, also known as fundamental forces (electromagnetism, gravitation, strong interaction, and weak interaction), in addition to being the