| Group → | 2 |
|---|---|
| ↓ Period | |
| 2 | 4 Be |
| 3 | 12 Mg |
| 4 | 20 Ca |
| 5 | 38 Sr |
| 6 | 56 Ba |
| 7 | 88 Ra |
The alkaline earth metals are a series of elements comprising Group 2 (IUPAC style) (Group IIA) of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). This specific group in the periodic table owes its name to their oxides that simply give basic alkaline solutions. These elements melt at such high temperature that they remain solids (“earths”) in fires. The alkaline earth metals provide a good example of group trends in properties in the periodic table, with well-characterized homologous behavior down the group. With the exception of Be and Mg, the metals have a distinguishable flame color, orange-red for Ca, magenta-red for Sr, green for Ba and crimson red for Ra.
Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:
| Z | Element | No. of electrons/shell |
|---|---|---|
| 4 | Beryllium | 2, 2 |
| 12 | Magnesium | 2, 8, 2 |
| 20 | Calcium | 2, 8, 8, 2 |
| 38 | Strontium | 2, 8, 18, 8, 2 |
| 56 | Barium | 2, 8, 18, 18, 8, 2 |
| 88 | Radium | 2, 8, 18, 32, 18, 8, 2 |
The alkaline earth metals are silver colored, soft metals, which react readily with halogens to form ionic salts, and with water, though not as rapidly as the alkali metals, to form strong alkaline (basic) hydroxides. For example, where sodium and potassium react with water at room temperature, magnesium reacts only with steam and calcium with hot water:
Beryllium is an exception: It does not react with water or steam, and its halides are covalent.
All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions.
The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These oxides are basic (alkaline) when combined with water. "Earth" is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating—properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths.
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Just like their names, they do not differ completely. The main difference is the electron configuration, which is ns2 for alkaline earth metals and ns1 for alkali metals. For the alkaline earth metals, there are two electrons that are available to for a metallic bond, and the nucleus contains an additional positive charge. Also, the elements of group 2A (alkaline earth) have much higher melting points and boiling points compared to those of group 1A (alkali metals). The alkali also have a softer and more lighweight figure whereas the alkaline earth metals are much harder and denser.
The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the alkali metals. The second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains much higher ionization energy, they still form an ionic compound with 2+ cations. Beryllium, however, behaves differently. This is because in order to remove two electrons from this particular atom, it requires significantly more energy. It never forms Be2+ and its bonds are polar covalent.
As mentioned earlier, Be is “special”; it behaves differently. If the Be2+ ion did exist, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since Be has a high charge density. All compounds that include Be have a covalent bond. Even the compound BeF2 , which is the most ionic Be compound, has a low melting point and a low electrical conductivity when melted.
Reactions:
Note: E = elements that act as reducing agents
1. The metals reduce halogens to form ionic halides: E(s) + X2 → EX2 (s) where X = F, Cl, Br or I
2. The metals reduce O2 to form the oxides: 2E(s) + O2 → 2EO(s)
3. The larger metals react with water to produce hydrogen gas: E(s) + 2H2O(l) → E2+(aq) + 2OH-aq + H2 (g) where E = Ca, Sr or Ba
Compounds
1. Alkylmagnesium halides (RMgX where R = hydrocarbon group and X = halogen). They are used to synthetise organic compounds.
Here’s an example: 3RMgCl + SnCl4 → 3MgCl2 + R3SnCl
2. Magnesium oxide (MgO). It is used as a material to refract furnace brick and wire insulation (melting point of 2852°C).
3. Calcium carbonate (CaCO3). It’s mainly used in the construction industry. It’s responsible for making limestone, marble, chalk, and coral.
| Alkaline earth metals | Atomic numbers in black indicate solids | Solid borders indicate primordial elements (older than the Earth) | Dashed borders indicate natural radioactive elements with no isotopes older than the Earth |
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ALKALINE EARTHS. The so-called alkaline earth-metals are the elements beryllium, magnesium, calcium, strontium and barium. By the early chemists, the term earth was used to denote those non-metallic substances which were insoluble in water and were unaffected by strong heating; and as some of these substances (e.g. lime) were found to be very similar in properties to those of the alkalis, they were called alkaline earths. The alkaline earths were assumed to be elements until 1807, when Sir H. Davy showed that they were oxides of various metals. The metals comprising this group are never found in the uncombined condition, but occur most often in the form of carbonates and sulphates; they form oxides of the type RO, and in the case of calcium, strontium and barium, of the type R02. The oxides of type RO are soluble in water, the solution possessing a strongly alkaline reaction and rapidly absorbing carbon dioxide on exposure; they are basic in character and dissolve readily in acids with the formation of the corresponding salts. As the atomic weight of the element increases, it is found that the solubility of the sulphates in water decreases.
Beryllium to a certain extent stands alone in many of its chemical' properties, resembling to some extent the metal aluminium. Beryllium and magnesium are permanent in dry air; calcium, strontium and barium, however, oxidize rapidly on exposure. The salts of all the metals of this group usually crystallize well, the chlorides and nitrates dissolve readily in water, whilst the carbonates, phosphates and sulphates are either very sparingly soluble or are insoluble in water.
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