What is the history of elements? Information about the elements in periodic table, isotopes, electron configuration, hybridization, reactivity of elements.
ELEMENT, a substance that cannot be decomposed or resolved into simpler substances by ordinary chemical means. It may also be defined as a substance that must gain mass on undergoing chemical reaction.
There are 105 known elements. Of these, 88 are found in nature in measurable quantities. The rest have been produced synthetically in laboratories. The elements may be combined into about four million compounds. About 95% contain carbon and are called organic compounds.
Elements are made up of atoms consisting of a very dense center, the nucleus, surrounded at relatively vast distances by electrons. The nucleus contains protons, which are particles with positive electrical charges, and neutrons, which are electrically neutral. Electrons have a negative charge. The mass of a proton and of a neutron is approximately the same, and the mass of each is about 1,840 times that of an electron. Protons and neutrons are held in th1 nucleus by other fundamental particles, such as mesons.
All atoms of an element contain the same number of positively charged protons in the nucleus and negatively charged electrons outside the nucleus. However, not all atoms of an element contain the same number of neutrons.
History. The ancient Greeks held a variety of opinions about the elemental composition of matter. Anaximenes of Miletus (6th century b. c. ) believed that water was the only element, while Empedocles of Agrigentum (5th century B.C.) and Aristotle (4th century b.c.) believed that there were four elements—earth, fire, air, and water. These elements were thought to be related to one another through four qualities—coldness, wetness, hotness, and dryness, which also gave the elements their characteristics. It was thought that one element could be transformed to another by changing the qualities.
Many variations of Aristotle’s theories were devised through the Middle Ages. In the 1500’s, for example, Paracelsus added three principles-sulfur, mercury, and salt—to the four elements. He felt that these principles had to be in balance in the body for good health.
Nine elements were known and used in the ancient world: carbon, copper, gold, iron, lead, mercury, silver, sulfur, and tin. The alchemists separated other elements from their compounds, and before 1600 antimony, arsenic, bismuth, and zinc were identified. Phosphorus was isolated during the 1600’s, while a number of elements, including hydrogen, nitrogen, and oxygen, were isolated in the 1700’s. About 30 elements were known by 1800, and over 80 by 1900.
The modern usage of the word “element” dates from Bobert Boyle in 1661, who described elements as “primitive and simple, or perfectly unmingled bodies; which not being made up of any other bodies, or of one another, are the ingredients of which all those called perfectly mixed bodies are immediately compounded, and into which they are ultimately resolved.”
In order to identify a substance as an element and classify it properly, it is first necessary to be able to separate the element from its compounds. For example, in his list of 33 elements published posthumously early in the 1800’s, the French chemist Antoine Lavoisier included 6 compounds. With the techniques available at the time these compounds could not be broken down into simpler substances. Second, it is necessary to have reliable atomic weights. By the mid-1800’s this was possible through the work of chemists such as Gay-Lussac, Berzelius, Faraday, Dumas, and Prout. Finally, sufficient numbers of elements had to be known so that the relationships between them could be clearly defined. About 60 elements were known in the mid-1800’s, enough to construct a system of classification.
The first real attempt to classify the elements was made about 1826 by the German chemist Johann Wolfgang Dobereiner. With his observation of triads, Dobereiner showed that in certain trios of elements—such as calcium, strontium, and barium—some of the properties of the middle element were predictable in terms of the properties of the other two elements. The elements in each triad were arranged by atomic weight.
In John Newlands’ law of octaves (1863) the elements were ordered according to atomic weight in eight columns of seven elements each. The eight elements of the horizontal rows showed repetitive properties. For example, in Newlands’ table, elements number 8 (fluorine), 15 (chlorine ), and 29 (iodine) are similar, while 22 (cobalt and nickel), 36 (palladium), and 50 (platinum and iridium) are similar. These elements differ by multiples of seven.
In 1869 the Russian chemist Dmitri I. Mendeleyev published his periodic law and table of elements. This table listed elements generally in order of increasing atomic weight. It also separated hydrogen from the other elements by placing it in its own period, and it left blank spaces for undiscovered elements. The table was published with the elements in horizontal rows. Each horizontal row is known as a period, or series; each vertical row is known as a family, or group. The families exhibited similar properties, so much so that Mendeleyev was able to predict the properties of the then undiscovered elements gallium, scandium, and germanium.
Two elements were not placed in the positions indicated by their atomic weights. Mendeleyev insisted that tellurium (atomic weight 127.60) should precede iodine (126.90) and that cobalt (58.93) should precede nickel (58.71). The order in these two cases was based on the properties of the elements. Mendeleyev’s table was based on the fact that the properties of the elements and their compounds are periodic functions of their atomic weights. The irregularities that he recognized were not explained until the 20th century.
In 1914 the English physicist Henry G. J. Moseley bombarded metal targets with electrons. He measured the wavelengths of the X rays produced and discovered a simple, linear, relationship between the square root of the reciprocal of the wavelength and Mendeleyev’s order of the elements. Moseley concluded that this regular increase must be due to the charge on the nucleus, and that it results from the successive addition of one proton to the nucleus in the progression from one element to the next through the periodic table. The number of protons in an atomic nucleus is called the atomic number.
When the elements are arranged in order of increasing atomic number, it is shown that Mendeleyev was correct in instances where he did not follow atomic weight.
Isotopes. Although each element has a characteristic number of protons (the atomic number) in the nuclei of its atoms, the number of neutrons may vary. Thus the total number of protons and neutrons (the mass number) in the nucleus may vary. Each of the variants of an element—identical in atomic number but different in mass number—is known as an isotope. For example, tin has 10 naturally occurring isotopes, each with 50 protons in its atoms but with different numbers of neutrons. Of the tin found in nature, 6.01« has 74 neutrons; 4.74«, 72; 32.75%, 70; 8.68%, 69; 23.84%, 68; 7.67%, 67; 14.28%, 66; 0.35%, 65; 0.68%, 64; and 1.01%, 62.
An isotope may be symbolized in various ways. For example, the tin isotope with mass number 120 may be written tin-120, Sn120, or 120Sn. The mass of an atom, measured in atomic mass units (1 atomic mass unit = 1/12 the mass of carbon-12), is nearly the same as its mass number. Tin-120, for example, has an atomic mass of 119.93. The atomic weight of an element is the average atomic mass of atoms occurring in nature, weighted according to their relative abundance. The atomic weight of tin is 118.69.
The chemical properties of isotopes are alike, but the physical properties that depend on mass differ. There are 269 stable isotopes that occur naturally and over 900 radioactive isotopes, many of which have been made in the laboratory.
Electron Configuration. Atomic electrons have certain severe restrictions in regard to the volume of space each may occupy. More exactly, the restrictions are on the probability of an electron occupying a certain volume. These restrictions are specified by quantum theory, which describes each electron in an atom by a set of four quantum numbers. According to Pauli’s exclusion principle, no two electrons in any single atom may have the same set of quantum numbers.
The principal quantum number, n, describes in a general way the size of the atom or its radius. The principal quantum numbers are numbered 1 through 7, with number 1 describing the energy level closest to the nucleus, and number 7 farthest from the nucleus. The electrons in an atom occur at seven definite, discrete energy levels, which may be loosely imagined as being distances from the nucleus. Electrons cannot occupy energy levels between these definite values. Inasmuch as energy is needed to remove a negative electron from an energy level closer to the positive nucleus to an energy level farther removed, it follows that the closer an electron is to the nucleus (the smaller its principal quantum number), the less energy it has.
The secondary quantum number, I, describes the shape of the volume of space in which the electron moves a certain percentage of the time, say 90%. The electron configurations identified by the primary and secondary quantum numbers have also been described in terms of shells and subshells surrounding the atomic nucleus. The shells, corresponding to the primary quantum numbers, are represented by the letters K, L, M, N, O, P, and Q. The subshells, corresponding to the secondary quantum numbers, are represented by the spectroscopic designations, s, p, d, and f. This alternate terminology, used for many years by spectroscopists, arises from the character of the lines produced in a spectrogram. An electron for which 1 = 0 produces a sharp line and is thus designated as an s electron; 1=1 produces a principle line; 1 = 2 produces a diffuse line; and I = 3 produces a fundamental line. Thus an electron in the second energy level (n = 2), whose secondary quantum number, I, is 1, can be written as 2p. (The spectroscopic designation s should not be confused with the spin quantum number, s. They are not related.)
The magnetic quantum number, m, describes the orientation in space of an electron with given principal and secondary quantum numbers. This orientation is made with respect to an external magnetic field, providing an arbitrary reference axis.
The spin quantum number, s, describes whether the spin of an electron around its own axis is clockwise or counterclockwise with reference to an external standard.
When electrons have identical quantum numbers except for spin, they are said to occupy the same orbital. An orbital may be regarded as a volume of space in which electrons move, and it may have one or two electrons but no more.
The restrictions upon the principle quantum numbers require that the first energy level has a maximum of 2 electrons, both s (one orbital); the second energy level has a maximum of 8 electrons (2s—one orbital, and 6p—three orbitals); the third energy level has a maximum of 18 electrons (2s, 6p, and lOd—five orbitals); the fourth energy level has a maximum of 32 electrons (2s, 6p, 10d, and 14/—seven orbitals). The fifth energy level has a theoretical maximum of 50 electrons, while the sixth may have 72 and the seventh, 98. Because only 105 elements are known at present, these maximum numbers of electrons are not found.
Hybridization. The magnetic quantum numbers, as well as the phenomenon known as hybridization, establish the three-dimensional structure of an atom in its compounds. Hybridization results from the tendency of electrons with somewhat different energies in different s, p, d, or f subshells to adopt equivalent energy positions. For example, carbon with its atomic number of 6 should have two s electrons in its first energy level and two s and two p electrons in its second energy level. However, the electrons in the second energy level hybridize so that the energy content of each is intermediate. This is sp3 (one s and three p) hybridization. Because they are all negatively charged, the electrons will repel one another, and the necessary geometrical structure in this case must be a tetrahedron. The angle from the center of a tetrahedron to two of its apexes is 109° 28′. Because of the equivalence of the four sp3 electrons in carbon, this angle is found in the carbon atom, in the diamond, which is pure carbon, and in symmetrical compounds of carbon, such as methane, CH4.
Reactivity of Elements. The reactivity of an element depends on a number of factors, including the electron complement in the outer shell (valence electrons) and the size of the atom. The electronegativity scale, which was devised separately by Linus Pauling and Robert Mulliken, rates the elements according to their tendency to gain or lose electrons. The scale assigns a number of 4.0 to fluorine and 0.7 to cesium. The closer an element is to 4.0, the greater its tendency to gain electrons, while the farther removed it is, the greater its tendency to lose electrons.
If the difference in electronegativity between two atoms of different elements in a compound is large, the compound will be ionic, as in sodium chloride, NaCl. In this compound an electron is removed from sodium, which forms a positively charged sodium ion, and is taken up by chlorine, which forms a negatively charged chloride ion. If the difference in electronegativity between the atom is small—as in methane, CH4—there will be a sharing of electrons and a covalent bond formed. The degree of covalency can be calculated from the difference in electronegativity between the constituent atoms.
If a large enough positive potential is applied to atoms, it will pull off one or more electrons, depending on its magnitude. Because positively charged ions result when electrons are pulled off, these potentials are called ionization potentials. If the electron is pulled off easily, the ionization potential is small. A table of ionization potentials gives data about reactivity and valences of the elements.
Elementary States. Elements may be isolated in each of the three states of matter. The state depends on the temperature and pressure.
The noble gases are helium, neon, argon, krypton, and xenon. Because their outermost electron shell is completely filled, they do not normally combine with other elements or even with like atoms. Therefore, there is only one atom per molecule of these gases. Most other gaseous elements contain two atoms per molecule —for example, hydrogen is H» and chlorine is Cl2. Some elements that are solid at room temperature vaporize under certain conditions to form gaseous molecules containing varying numbers of atoms. For example, phosphorus may have four atoms per molecule, P2, or at higher temperatures, two atoms per molecule P2. Sulfur vapor, depending on the temperature, may be S3, S6, or S8.
A few of the elements, such as bromine, Br2, mercury, Hg, and perhaps gallium, Ga, are liquid at room temperature.
The solid elements generally form crystalline structures of various types. The shape of the crystal is directly dependent on the arrangement of the atoms of the element.
Elements may exist in two or more different forms, such as oxygen, O2, and ozone, O3; gray and red selenium; and white and red phosphorus. The phenomenon of different forms of the same element is known as allotropy. The term polymorphism is also used for different crystalline forms of an element.