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Gaffer Variety:
Elements 3 SN2 005:
By Willie Gaffer:
This is a good enough place to make a very important point. Chemistry is not magic or necromancy. It just seems that way because those who present it do not explain it properly. None of the five books I have tried to use explains the periodic table, the structure of the atom, and valence in a forthright way. To be sure, the information is usually presented, but it is not properly organized. It is scattered hither and yon thought the text. If the ideas are properly ordered, they can be laid out in just a few pages so a child could understand them.
With that in mind, let us discuss valence. This is the most
important concept in chemistry.
So, chemical compounds are formed by the joining of two or more atoms. We have a stable compound when the total energy of the combination has lower energy than the separated atoms. The bound state implies a net attractive force between the atoms, which creates a chemical bond. There are two main ways in which bonding can occur. First, a covalent bond is one in which one or more pairs of electrons are shared by two atoms. Second, an ionic bond is one in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions, which attract each other. There are other forms of bonds that can occur but we need not consider those immediately.
Now we can discuss the series arrangements of the periodic table. The sets or series of elements do not correspond exactly to the column or group designations. These series designation will vary depending upon which reference we are using. This is a large problem we have in chemistry and is one reason why so many people think of it as necromancy. When the professors and chemists cannot even agree amongst themselves on what to call something, the critics are somewhat justified.
By series we mean the groupings of the kinds of elements we have. For example group 1A, excepting hydrogen, is called the alkali metals series. They are called alkali metals because they form alkalis when they combine with other elements. Alkalis are strong bases capable of neutralizing acids. Sodium and potassium are the most common of these. The other alkali metals are much more rare. Francium, a natural radioactive isotope, is very rare and was not discovered until 1939.
The alkali metals are extremely reactive and combine readily with most of the substances found in the atmosphere. Because of their high reactivity, the alkali metals are never found as free metals in their natural state. They generally are found combined with other elements in the form of simple or complex compounds. Salt (sodium chloride, NaCl), is one combination we are all familiar with. Although most of us think of it as a useful seasoning, it has enormous commercial value and entire industries are devoted to its mining and use.
The six elements of group 2A are called the alkaline earth metals. The alkaline-earth metals are extremely electropositive. That means, like the alkali metals of group IA, their atoms easily lose electrons to become positive ions called cations. Most of their typical compounds are ionic, that is salts in which the metal occurs as uniformly divalent cations. For now, a divalent bond is one in which the valence is 2. The salts are colorless unless they include a colored anion (negative ion). Typical alkaline-earth compounds are calcium chloride (CaCl2) and calcium oxide (CaO). These are different from the alkali metals, which form monovalent compounds like sodium chloride (NaCl) and sodium monoxide (Na2O). The oxides of the alkaline-earth metals are basic meaning alkaline, in contrast to acidic.
All the metals and their compounds find commercial application to some degree, especially magnesium alloys and a variety of calcium compounds. Magnesium and calcium, particularly the latter, are abundant in nature and play significant roles in geologic and biological processes. Radium is a rare element; all its isotopes are radioactive.
Now we have the rare earth elements. In the periodic table of the elements the rare-earth elements comprise three members of Group IIIb and all 14 members of one of the two series of elements that hang below the main table. This long series is known collectively as the lanthanide series because it directly follows lanthanum in a different form of the table. The rare-earth elements all have certain common features in the electronic structure of their atoms, which is the fundamental reason for their chemical similarity.
The chemistry of all the rare earths is very similar and changes only slightly in progressing along the lanthanide series. All of these elements form trivalent compounds, and in the crystal lattices (the regular arrangement of atoms in the solid forms) of such compounds, one rare-earth ion readily replaces another. The rare-earth metals when heated react strongly with nonmetallic elements to form very stable compounds. They are never found as the free metals in the Earth's crust. Pure minerals of individual rare earths do not exist in nature; all their minerals contain mixtures of the rare-earth elements.
Next are the actinide elements. This is the series of 15 consecutive chemical elements in the periodic table from actinium to lawrencium (atomic numbers 89–103). All but actinium hang below the table. As a group the actinides are significant largely because of their radioactivity. Although several members of the group, including uranium (the most familiar), occur naturally, most are man-made. Both uranium and plutonium have been used in atomic bombs for their explosive power and currently are being employed in nuclear plants for the production of electrical power. Although there are a number of exceptions to these generalities, for most of these elements, the concept of a series of nearly identical actinide elements is a useful guide for predicting their chemical and physical properties.
Of course, each actinide has its own particular atomic number, however the atoms of an element are capable of existing in a number of forms (isotopes), each of which has a different number of neutrons in its nucleus and hence a different atomic mass. Although isotopes of a given element behave alike chemically, they may have different stabilities in relation to radioactive decay, which is a property of the nucleus. No element beyond bismuth in the periodic classification has any stable isotopes.
The actinides are unusual in forming a series of 15 elements having no stable isotopes; every actinide isotope undergoes radioactive decay, and, as a result, only a few of the lighter, more stable members of the series (such as thorium and uranium) are found in nature. The half-life, or the precise time required for one-half of any amount of a particular isotope to disappear due to radioactive decay, is a measure of the stability of that isotope. The naturally occurring isotopes in the actinide series have long half-lives, of the order of billions of years.