Alkali Metals - Group IA

They are so called because they form alkalies (soluble bases). Sodium and potassium are the sixth and seventh most abundant of the elements, comprising, respectively, 2.6 and 2.4 percent of the Earth's crust.

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. The simpler compounds of the alkali metals are soluble in water and therefore are easily extracted and subjected to chemical operations for purposes of separation and purification. Minerals belonging to this class, such as halite (NaCl), sylvite (KCl), and carnallite (KCl × MgCl2 × 6H2O), although somewhat rare, are the most important commercial sources of the alkali metals.

The alkali metals have all of the physical properties generally associated with metals, including silver-like lustre, high ductility, and excellent conductivity of electricity and heat. Lithium is the lightest metallic element. The alkali metals are low melting, ranging from a high of 179° C for lithium to a low of 28.5° C for cesium. Alloys of alkali metals exist that melt as low as -78° C.

The alkali metals are extremely reactive and combine readily with most of the substances found in the atmosphere. (Only lithium, however, reacts with nitrogen.) The alkali metals all react vigorously, and often violently, with water, releasing hydrogen and forming strong caustic solutions. Most common non-metallic substances such as the halogens, halogen acids, sulphur, and phosphorus react with the alkali metals. The alkali metals themselves react with many organic compounds, particularly those containing an active hydrogen atom.

Alkali Earth Metals Group- II A

Prior to the 19th century, substances that were nonmetallic, insoluble in water, and unchanged by
fire were known as earths. Those earths, like lime, that resembled the alkalies (soda ash and potash) were designated alkaline earths. Alkaline earths were thus distinguished from the alkalies and from other earths, such as alumina and the rare earths. By the early 1800s it became clear that the earths, formerly considered to be elements, were in fact oxides, compounds of a metal and oxygen. The metals whose oxides make up the alkaline earths then came to be known as the alkaline-earth metals and have been classified in group II of the periodic table ever since Mendeleyev proposed his first table in 1869.

The alkaline-earth metals are extremely electropositive; that is, like the alkali metals, their atoms easily lose electrons to become positive ions (cations). The salts are colourless unless they include a coloured anion (negative ion).
The oxides of the alkaline-earth metals are basic. A fairly steady increase in electropositive character is observed in passing from beryllium, the lightest member of the group, to radium, the heaviest; as a result of this trend, beryllium oxide is only weakly basic and even shows acidic properties, whereas barium and radium oxide are strongly basic. The metals themselves are highly reactive reducing agents; that is, they readily give up electrons to other substances that are, in the process, reduced.

Halogens - Group VII A

They were given the name halogen from the Greek roots hal- (“salt”) and gen (“to produce”), because they all produce sodium salts of similar properties, of which sodium chloride, table salt, is the best known.

The free halogen elements are not found in nature because of their great reactivity. In combined form, fluorine is the most abundant of the halogens in the Earth's crust. The percentages of the halogens in the igneous rocks of the Earth's crust are 0.06 fluorine, 0.031 chlorine, 0.00016 bromine, and 0.00003 iodine. Astatine does not occur in nature because it consists only of short-lived radioactive isotopes.

The halogen elements show great resemblances to one another in their general chemical behaviour and in the properties of their compounds with other elements. There is, however, a progressive change in properties from fluorine through chlorine, bromine, and iodine to astatine—the difference between two successive elements being most pronounced with fluorine and chlorine. Fluorine is the most reactive of the halogens and, in fact, of all elements, and it has certain other properties that set it apart (see below General properties of the group).

Noble Gases - Group 0

These are so called because they have completed their octets and have thus achieved stability. They rarely react with any element and whenever they do so it is at high temperatures. The term “noble” alludes to the extraordinarily limited reactivity of the gases. They also are sometimes called inert gases for the same reason. It was, however, discovered in 1962 that the heavier noble gases—krypton, xenon, and radon—can form chemical compounds with fluorine, the strongest electron attracting of all elements. The outer electrons of the atoms of these three gases, screened from the nucleus by intervening electrons, are held less firmly and can be removed.

After hydrogen, helium is the most plentiful element in the universe, comprising almost 25 percent of its total mass. Under ordinary conditions, the noble gases are colourless, odourless, and non-flammable. The noble gases absorb and give off electromagnetic radiation in a much less complex manner than do other substances. This absorption and emission behaviour is exploited in the use of the gases (with the exception of highly radioactive radon) in fluorescent lighting devices and discharge lamps. If any of the noble gases is confined at low pressure in a glass tube and an electrical charge is passed through it, the gas glows. The noble gases also have very low boiling points and melting points, which make them useful as refrigerants in the study of matter at extremely low temperatures.