Alkaline Earth Metals - How it works



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Defining a Family

The expression "families of elements" refers to groups of elements on the periodic table that share certain characteristics. These include (in addition to the alkaline earth metals and the alkali metals) the transition metals, halogens, noble gases, lanthanides, and actinides. (All of these are covered in separate essays within this book.) In addition, there are several larger categories with regard to shared traits that often cross family lines; thus all elements are classified either as metals, metalloids, and nonmetals. (These are also discussed in separate essays, which include reference to "orphans," or elements that do not belong to one of the families mentioned above.)

These groupings, both in terms of family and the broader divisions, relate both to external, observable characteristics, as well as to behaviors on the part of electrons in the elements' atomic structures. For instance, metals, which comprise the vast majority of elements on the periodic table, tend to be shiny, hard, and malleable (that is, they can bend without breaking.) Many of them melt at fairly high temperatures, and virtually all of them vaporize (become gases) at high temperatures. Metals also form ionic bonds, the tightest form of chemical bonding.

ELECTRON CONFIGURATIONS OF THE ALKALINE EARTH METALS.

Where families are concerned, there are certain observable properties that led chemists in the past to group the alkaline earth metals together. These properties will be discussed with regard to the alkaline earth metals, but another point should be stressed in relation to the division of elements into families. With the advances in understanding that followed the discovery of the electron in 1897, along with the development of quantum theory in the early twentieth century, chemists developed a more fundamental definition of family in terms of electron configuration.

As noted, the alkaline earth metal family occupies the second group, or column, in the periodic table. All elements in a particular group, regardless of their apparent differences, have a common pattern in the configuration of their

DURING WORLD WAR II, MAGNESIUM WAS HEAVILY USED IN AIRCRAFT COMPONENTS. IN THIS 1941 PHOTO, WORKERS POUR MOLTEN MAGNESIUM INTO A CAST AT THE WRIGHT AERONAUTICAL CORPORATION. (Bettmann/Corbis. Reproduced by permission.)
D URING W ORLD W AR II, MAGNESIUM WAS HEAVILY USED IN AIRCRAFT COMPONENTS . I N THIS 1941 PHOTO , WORKERS POUR MOLTEN MAGNESIUM INTO A CAST AT THE W RIGHT A ERONAUTICAL C ORPORATION . (
Bettmann/Corbis
. Reproduced by permission.)
valence electrons—the electrons at the "outside" of the atom, involved in chemical bonding. (By contrast, the core electrons, which occupy lower regions of energy within the atom, play no role in the bonding of elements.)

All members of the alkaline earth metal family have a valence electron configuration of s 2 . This means that two electrons are involved in chemical bonding, and that these electrons move through an orbital, or range of probabilities, roughly corresponding to a sphere. The s orbital pattern corresponds to the first of several sub-levels within a principal energy level.

Whatever the number of the principal energy level which corresponds to the period, or row, on the periodic table the atom has the same number of sublevels. Thus beryllium, on Period 2, has two principal energy levels, and its valence electrons are in sublevel 2 s 2 . At the other end of the group is radium, on Period 7. Though radium is far more complex than beryllium, with seven energy levels instead of two, nonetheless it has the same valence electron configuration, only on a higher energy level: 7 s 2 .

HELIUM AND THE ALKALINE EARTH METALS.

If one studies the valence electron configurations of elements on the periodic table, one notices an amazing symmetry and order. All members of a group, though their principal energy levels differ, share characteristics in their valence shell patterns. Furthermore, for the eight groups numbered in the North American version of the periodic table, the group number corresponds to the number of valence electrons.

There is only one exception: helium, with a valence electron configuration of 1 s 2 , is normally placed in Group 8 with the noble gases. Based on that s 2 configuration, it might seem logical to place helium atop beryllium in the alkaline earth metals family; but there are several reasons why this is not done. First of all, helium is obviously not a metal. More importantly, helium behaves in a manner quite different from that of the alkaline earth metals.

Whereas helium, like the rest of the noble gases, is highly resistant to chemical reactions and bonding, alkaline earth metals are known for their high reactivity—that is, a tendency for bonds between atoms or molecules to be made or broken so that materials are transformed. (A similar relationship exists in Group 1, which includes hydrogen and the alkali metals. All have the same valence configuration, but hydrogen is never included as a member of the alkali metals family.)

Characteristics of the Alkaline Earth Metals

Like the alkali metals, the alkaline earth metals have the properties of a base, as opposed to an acid. The alkaline earth metals are shiny, and most are white or silvery in color. Like their "cousins" in the alkali metal family, they glow with characteristic colors when heated. Calcium glows orange, strontium a very bright red, and barium an apple green. Physically they are soft, though not as soft as the alkali metals, many of which can be cut with a knife.

Yet another similarity the alkaline earth metals have with the alkali metals is the fact that four of the them—magnesium, calcium, strontium, and barium—were either identified or isolated in the first decade of the nineteenth century by English chemist Sir Humphry Davy (1778-1829). Around the same time, Davy also isolated sodium and potassium from the alkali metal family.

REACTIVITY.

The alkaline earth metals are less reactive than the alkali metals, but like the alkali metals they are much more reactive than most elements. Again like their "cousins," they react with water to produce hydrogen gas and the metal hydroxide, though their reactions are less pronounced than those of the alkali metals. Magnesium metal in its pure form is combustible, and when exposed to air, it burns with an intense white light, combining with the oxygen to produce magnesium oxide. Likewise calcium, strontium, and barium react with oxygen to form oxides.

Due to their high levels of reactivity, the alkaline earth metals rarely appear by themselves in nature; rather, they are typically found with other elements in compound form, often as carbonates or sulfates. This, again, is another similarity with the alkali metals. But whereas the alkali metals tend to form 1+ cations (positively charged atoms), the alkaline earth metals form 2+ cations—that is, cations with a positive charge of 2.

BOILING AND MELTING POINTS.

One way that the alkaline earth metals are distinguished from the alkali metals is with regard to melting and boiling points—those temperatures, respectively, at which a solid metal turns into a liquid, and a liquid metal into a vapor. For the alkali metals, the temperatures of the boiling and melting points decrease with an increase in atomic number. The pattern is not so clear, however, for the alkaline earth metals.

The highest melting and boiling points are for beryllium, which indeed has the lowest atomic number. It melts at 2,348.6°F (1,287°C), and boils at 4,789.8°F (2,471°C). These figures are much higher than for lithium, the alkali metal on the same period as beryllium, which melts at 356.9°F (180.5°C) and boils at 2,457°F (1,347°C).

Magnesium, the second alkali earth metal, melts at 1,202°F (650°C), and boils at 1,994°F (1,090°C)—significantly lower figures than for beryllium. However, the melting and boiling points are higher for calcium, third of the alkaline earth metals, with figures of 1,547.6°F (842°C) and 2,703.2°F (1,484°C) respectively. Melting and boiling temperatures steadily decrease as energy levels rise through strontium, barium, and radium, yet these temperatures are never lower than for magnesium.

ABUNDANCE.

Of the alkaline earth metals, calcium is the most abundant. It ranks fifth among elements in Earth's crust, accounting for 3.39% of the elemental mass. It is also fifth most abundant in the human body, with a share

MANY ALKALINE EARTH METALS ARE USED IN THE PRODUCTION OF FIREWORKS. (Bill Ross/Corbis. Reproduced by permission.)
M ANY ALKALINE EARTH METALS ARE USED IN THE PRODUCTION OF FIREWORKS . (
Bill Ross/Corbis
. Reproduced by permission.)
of 1.4%. Magnesium, which makes up 1.93% of Earth's crust, is the eighth most abundant element on Earth. It ranks seventh in the humanbody, accounting for 0.50% of the body's mass.

Barium ranks seventeenth among elementsin Earth's crust, though it accounts for only0.04% of the elemental mass. Neither it nor the other three alkali metals appear in the body insignificant quantities: indeed, barium and beryllium are poisonous, and radium is so radioactivethat exposure to it can be extremely harmful.

Within Earth's crust, strontium is present in quantities of 360 parts per million (ppm), which in fact is rather abundant compared to a number of elements. In the ocean, its presence is about 8 ppm. By contrast, the abundance of beryllium in Earth's crust is measured in parts per billion (ppb), and is estimated at 1,900 ppb. Vastly more rare is radium, which accounts for just 0.6 parts per trillion of Earth's crust—a fact that made its isolation by French-Polish physicist and chemist Marie Curie (1867-1934) all the more impressive.

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