Group 2 elements are reactive metals, though their behavior is moderated compared to their Group 1 neighbors. These highly metallic elements are never found in nature in their pure, elemental form because they readily combine with other substances. They are considered the second most reactive family of elements, after the alkali metals of Group 1. This chemical restlessness means pure samples of these metals, like strontium and barium, must be stored under a protective layer of liquid paraffin or oil to prevent immediate reaction with the atmosphere.
Defining the Alkaline Earth Metals
The elements of Group 2 are collectively known as the Alkaline Earth Metals, a name that reflects both their chemical properties and their natural occurrence. This group includes beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and the radioactive element radium (Ra). A defining feature of these metals is the presence of exactly two electrons in their outermost electron shell.
These two outer electrons drive their chemistry. The atoms are just two steps away from achieving the highly stable electron configuration of a noble gas. To attain this stability, the metals readily shed both electrons to form a positively charged divalent cation with a charge of +2. This tendency to lose electrons is a characteristic metallic property that dictates their behavior in chemical reactions.
The Driving Force Behind Group 2 Reactivity
The underlying reason for the high reactivity of Alkaline Earth Metals is their inherent drive toward a stable electron arrangement. Achieving the electron structure of the nearest noble gas requires them to lose their two valence electrons, gaining a full, stable outer shell.
The energy required to remove an electron from an atom is called the ionization energy. Group 2 elements have relatively low ionization energies for the first two electrons, making it easy to shed them and form the +2 ion. However, the energy cost to remove a third electron is dramatically higher, as it must be pulled from the now-stable inner shell, which is held much more tightly.
This sharp increase in the third ionization energy locks the element into forming only a +2 ion in almost all chemical compounds. The ease with which they lose the first two electrons justifies their classification as highly reactive metals. Because these elements readily lose electrons (a process called oxidation), they act as strong reducing agents in reactions.
How Reactivity Changes Across the Group
The reactivity of the Alkaline Earth Metals is not uniform; it increases noticeably as one moves down the group from Beryllium to Barium. This trend is directly linked to the changing size of the atoms down the column. Moving down the group, each successive element adds a new electron shell, causing the atomic radius to increase significantly.
As the atom gets larger, the two outermost valence electrons are farther away from the positive pull of the nucleus. This greater distance, combined with the shielding effect of the inner electron shells, weakens the attractive force on the valence electrons. Consequently, the ionization energy decreases down the group, meaning less energy is required to remove those two electrons.
This results in a gradient of reactivity, with Beryllium being the least reactive and Barium being the most reactive element in the group. Beryllium, for instance, exhibits some unique properties, such as forming compounds with more covalent character, and it does not react with water even when heated. Conversely, Barium reacts vigorously with most substances, demonstrating the increased ease of electron loss at the bottom of the group.
Common Chemical Interactions
The inherent reactivity of Group 2 metals is most clearly demonstrated in their common reactions with oxygen and water. When exposed to air, all Group 2 metals react to form metal oxides, which is why pure samples often appear dull due to a surface tarnish. This reaction is a reduction-oxidation process where the metal is oxidized from a zero to a +2 oxidation state. For example, magnesium metal reacts with oxygen to produce magnesium oxide.
Their reaction with water follows the reactivity trend down the group, producing hydrogen gas and a metal hydroxide, which makes the resulting solution alkaline. Magnesium reacts very slowly with cold water, but will react vigorously with steam to form magnesium oxide. Calcium, strontium, and barium react progressively more vigorously with cold water, with barium reacting quite rapidly.
Another visible manifestation of their chemical nature is the characteristic color they impart to a flame when heated. Calcium produces a brick-red color, strontium gives a crimson hue, and barium yields a green color. These characteristic flame tests are used analytically to confirm the presence of the metal ions.