BASIC CONCEPTS: THE PERIODIC TABLE, SYMBOLS AND NOTATION, AND COMMON QUANTITIES AND THEIR UNITS

SYMBOLS AND NOTATION

A. As previously mentioned, chemistry is based on the premise that all matter is composed of some combination of the 92 naturally occurring elements. [Although the periodic table lists 118 elements, the final 26 are man-made or artificially synthesized, radioactive, and increasingly unstable, that is, they have short half-lives. They are not addressed in this chapter, but are discussed in Chapter 5 of “Applied Chemistry for Environmental Engineering,” the companion volume to this book.] These elements are arranged in a particular order, known as the periodic chart or periodic table, in horizontal rows called periods, by increasing atomic number, and in vertical columns called groups or families, by similar electron arrangements or configurations, which give rise to similar chemical properties. Electron configurations are ordered arrangements of electrons based on specific housing rules. Energy levels or shells are quantized and exist from one to infinity. In the ground state of an atom, however, electrons can populate energy levels ranging from 1 to 7. Within each energy level or shell are subshells containing orbitals, the first four of which are labeled s, p, d, and f Each of the four orbitals has a distinct shape and population limit: 1 s orbital with a maximum of 2 electrons; 3 p orbitals with a maximum of 6 electrons; 5 d orbitals with a maximum of 10 electrons; and 7 f orbitals with a maximum of 14 electrons.

B. Members of the same family are called congeners, especially important in organizing organic compounds such as polychlorinated biphenyls (or PCBs) or organochlorine pesticides. Here, the PCB molecule may have two, four, six, or eight chlorine atoms; all are congeners of the same family. Of more immediate relevance is the fact that some rows of the periodic table have names, such as the actinide (row 6) or lanthanide (row 7) series, while columns have names such as alkali metals (column 1A or 1), alkaline earth metals (column 2A or 2), transition metals (columns 3 through 12), halogens (column 7A or 17), and noble (inert) gases (column 8A or 18). The pure numeral system of identifying columns is more modern and preferred over the number and letter system.

C. Each element is symbolized by either a single capital letter or a capital letter followed by a single lowercase letter. Thus, the number of elements from which a compound is made up can easily be determined by counting the number of capital letters.

D. The atomic number, Z, of an element is the number of protons in the nucleus of its atom. The atomic number, which is always an integer and uniquely identifies the element, is written as a left-hand subscript to the element symbol, for example, 6C.

E. The mass number, A, of an element is the sum of the number of protons and number of neutrons in the nucleus, collectively known as the number of nucleons. Protons and neutrons are assigned a mass number of one atomic mass unit (amu) each, based on the carbon-12 atom as the standard. Although an element can have only one atomic number, it may have more than one mass number. This fact gives rise to the phenomenon of isotopes. Isotopes are atoms of the same element with the same atomic number but with a different number of neutrons. An element may have several isotopes, one or more of which may be radioactive and hence unstable. For example, oxygen has three naturally occurring isotopes, all of which are stable, while carbon also has three, one of which is radioactive. The mathematical average of all isotopic mass numbers of an element, weighted by the percent abundance in nature of these isotopes, constitutes the element’s atomic mass. It is impossible to predict theoretically how many isotopes an element may have, or how many may be radioactive. The isotopes of a given element have almost identical chemical properties (i.e., reactivity) but different physical properties (i.e., density, melting point, etc.). The mass number is expressed as a left- hand superscript to the element symbol, for example, 12C.

F. In its elemental state, the electronic charge of an element is zero, that is, the number of electrons equals the number of protons. If an atom gains or loses electrons, it becomes charged. A charged atom is called an ion. A positively charged ion has lost one or more electrons and is called a cation, while a negatively charged ion has gained one or more electrons and is called an anion. The charge or oxidation state on the ion is expressed as a right-hand superscript to the element symbol, for example, Ca2+.

G. Elements in a given vertical column are referred to as a group or family. The reason for this is that they have similar electron configurations and thus prefer to gain or lose the same number of electrons. This, in turn, helps predict their reactivity with other elements. There are eight major groups or families. For example, elements in the first two columns are known as alkali metals (starting with lithium and ending with francium) and alkaline earth metals (starting with beryllium and ending with radium), respectively. They prefer to lose one and two electrons from their outermost energy levels, respectively. They form cations, just as other metals do. Their ions are electron deficient; hence the number of protons outnumbers the number of electrons. Alkali metals are said to have a valence or oxidation state of +1, or are univalent, while alkaline earth metals are said to have a valence or oxidation state of +2, or are divalent. Meanwhile, elements in the next to the last column, known as halogens, are nonmetals. They prefer to gain one electron and form an anion, and have a valence or oxidation state of -1. Both the tendency for metals to lose one or more electrons and the tendency for nonmetals to gain one or more electrons are explained by their electron configurations and energy stabilization rules, and account for the respective reactivities. There is a large array of metals in the middle of the periodic table known as the transition metals. Because of their complex electron configurations, they have multiple, energetically stable, positive valences or oxidation states and form a variety of interesting, industrially useful compounds. The seven elements that lie in between metals and nonmetals that form a descending step in the periodic table are known as metalloids. They have the ability to both gain and lose electrons, depending on their immediate chemical environment; their properties are often less predictable but equally interesting. Consult Table 1.1 in the next section to see which elements form cations and which form anions, and which may form both.

In general, oxides of nonmetals, when dissolved in water, are acidic. Oxides of metals are alkaline or basic. Thus, substances such as calcium oxide or magnesium oxide are slightly alkaline in the presence of water. Substances such as carbon dioxide or sulfur dioxide, on the other hand, are slightly acidic. Normal rainwater, for example, has a pH of about 5.6 (slightly acidic) due to the presence of carbon dioxide in the atmosphere. For a distribution of important elements in the environment, see Table 1.8 at the end of this chapter.

H. Free radicals and ions may occasionally be confused with each other, but they are quite different and have different stabilities. Ions are electrically charged atoms, having gained or lost one or more electrons, for reasons discussed in the preceding paragraph, (G). They are quite stable in aqueous salt solutions. A list of common cations and anions is provided in Table 1.1.

Free radicals, on the other hand, concern unpaired electrons in electrically neutral molecules. Stable molecules are made when two atoms join together and form a bond in which two electrons are shared. This is a single bond and is consistent with the duet and octet rules for stable molecules. As will be seen in Chapter 3 in organic chemistry, double and triple bonds are also possible, in which four and six electrons are shared, respectively. In addition to bonding pairs, molecules may also have lone pairs of electrons, which do not participate in the bonding process, but serve to fulfill the octet rule for valence electrons. They also exist as pairs. However, occasionally, a molecule may exist in a metastable state having resonance structures (discussed in Chapter 3), where an atom has an unpaired electron. Such a molecule is referred to as a free radical, is highly reactive, and has a very short lifetime. One example of a free radical is the hydroxyl radical *OH [not to be confused with the hydroxide anion OH-]. The equation below shows the reaction of hydroxyl radical with a methane molecule to produce a methyl radical, another free radical:

From a health perspective, free radicals are known to react with DNA in the human body, leading to its unraveling and breakdown in coding operations and proper replication. Consumption of antioxidants, such as L-ascorbic acid (vitamin C), which preferentially react with free radicals, is believed to reduce this problem.

Example 1.1

Uranium, the element with atomic number Z = 92, has three naturally occurring isotopes: U-234, U-235, and U-238. Only U-235 is fissionable. When it combines with a halogen to form a compound, it commonly forms a cation with a +6 charge.

A. Using proper notation, write the symbol of the U-235 isotope, indicating its electronic charge (also known as the valence or oxidation state).

B. Determine the number of protons, neutrons, and electrons in one atom of this isotope.

Solution

A. The symbol, written with proper notation, is

B. The atom has 92 protons (note left-hand subscript).

It must also have 92 electrons, understood, in its neutral (uncharged) atom. However, since this atom has a charge of +6 (note right-hand superscript), it has lost six electrons to become positively charged. Hence, it now has only 86 electrons.

To compute the number of neutrons in this atom, subtract the atomic number (left-hand subscript) from the mass number (left- hand superscript): 235 - 92 = 143 neutrons.

 
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