Chemistry of the Group 2 Elements - Be, Mg, Ca, Sr, Ba, Ra Content from the guide to life, the universe and everything

Chemistry of the Group 2 Elements - Be, Mg, Ca, Sr, Ba, Ra

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Overview of the Periodic Table
Hydrogen | Group 1 - Alkali Metals | Group 2 - Alkaline Earth Metals | Group 12













Alkaline earth metals can be found in various minerals in the Earth's crust. Beryllium is found in the mineral beryl, Be3Al2(SiO3)6, which must be handled with care as beryllium compounds are highly toxic. Calcium can be widely found in limestone, CaCO3, and in dolomite along with magnesium, CaCO3.MgCO3. Magnesium can also be found in carnallite, KCl.MgCl2.6H2O. Strontium and barium can be found in the less abundant minerals, strontianite (SrCO3) and barytes (BaSO4). These elements can also be found as soluble salts in sea water. All of the isotopes of radium are radioactive and occur in the decay pathways for uranium-238.

These elements have smaller atomic radii than the group 1 elements due to the greater nuclear charge and with two valence electrons the material have much higher melting and boiling points. They also have high ionisation energies but this is made up for by large lattice and hydration enthaplies. The chemistry of these elements is somewhat like that of the group 12 elements, zinc, cadmium and mercury.

Preparation, Properties and Uses of the Elements

Metallic beryllium can be prepared by the reduction of BeF2 by magnesium or by the Ca or Mg reduction or electrolysis of BeCl2. It is the least reactive element in its group and is unreactive with water and air unless heated to very high temperature. Due to its very small atomic radius and high ionisation energies, lattice and hydration enthalpies aren't large enough to achieve the charge separation necessary to form simple Be2+ions in its compounds, and so its chemistry is largely covalent. If it is powdered, it can be ignited, reacting with air to give a mixture of BeO and Be3N2. It is unreactive with cold nitric acid, but will react with non-complexing acids1 to produce the tetraaqua ion [Be(H2O)4]2+. Beryllium is one of the lightest known materials and has a very high melting point for a light metal. It is used to make high strength alloys of copper and nickel and is used to make the windows of X-ray tubes.

Magnesium is made on large scales by two main processes:

  • Electrolysis of mixtures of metal halide salts

  • The reduction of calcined dolomite by ferrosilicon alloy

CaCO3.MgCO3 → CaO.MgO + 2 CO2
2 (CaO.MgO) + FeSi → 2 Mg + Ca2SiO4 + Fe

The reactivity of Mg is between that of Be and the rest of the alkaline earth elements. The Mg2+ ion is very small and highly polarizing, so much of its chemistry can also be covalent. It burns in air and reacts with steam to form MgO and will react with dilute acids. It will react with ammonia and nitrogen gas at high temperatures to form the nitride Mg3N2. It is widely used in construction in light weight alloys and is also used to make Grignard reagents, described later in this entry. It is also a very important biological metal - it is present in the pigment 'chlorophyll', which absorbs light in the photosynthetic processes in plants, and is also the ion that neutralises the negative charge carried by the phosphate groups in DNA.

Calcium, strontium and barium are readily made from the reduction of their halides using sodium and are produced on much smaller scales than beryllium and magnesium. Calcium is used as a reducing agent to isolates lanthanide and actinide elements from their halides and also to make CaH2, a powerful reducing agent and a drying agent for chlorinated solvents.

Table 1: Some properties of the group 2 elements.
ElementM2+ ionic radius (Å)2Melting point (°C)Boiling point (°C)First and second ionisation enthalpies (kJ mol-1)
Be0.5912872500899, 1757
Mg0.866491105738, 1450
Ca1.148391494590, 1145
Sr1.327681381549, 1064
Ba1.497271850503, 965
Ra1.627001700509, 975


As has been mentioned, the chemistry of beryllium is highly covalent in nature. The beryllium atom has two valence electron and can therefore form two covalent bonds. There are, however, four available atomic orbitals in the valence shall of beryllium, so molecules of the type BeX2, which are said to coordinatively unsaturated, can associate to formed an extended covalently bonded network. Simple BeX2 type monomeric molecules only exist in the gas phase.

Beryllium will form polymeric compounds with the formula (BeX2)n, where X acts as a bridging group between adjacent Be atoms. Compounds of this structure occur where X is H, F, Cl and CH3. Each Be atom is therefore able to form four equivalent covalent bonds which are arranged in a tetrahedral geometry (figure 1, drawn in 2D). Compounds can be formed which have the structure (M+)2(Be4Cl10)2-, where the (Be4Cl10)2- resembles a stretch of the (BeCl2)n polymer chain. (BeCl2)n is best made through the reaction of elemental beryllium with gaseous Cl2 with heating. (BeF2)n can be made by thermally decomposing NH4(BeF4) which in turn is made by dissolving BeO in an aqueous solution containing fluoride ions.

Figure 1: Structure of polymeric (BeX2)n. X X X \ / \ / \ / \ / … Be Be Be Be … / \ / \ / \ / \  X X X

Similar compounds can also be made where X is an alkoxide RO, where R os some organic group. For example, the compound (Be(OCH3)2)n is also polymeric. However if R is much bigger than CH3, such as C(CH3)3 (written in shorthand as But), the bridging groups clash together preventing an extended structure, allowing only the trimer (Be(OBut)2)3 to form. In this compound, the central Be atom forms four bonds in a tetrahedral arrangement whilst the outer two only form three bonds each (figure 2).

Figure 2: Structure of (Be(OBut)2)3 But But O O / \ / \  ButO-Be Be Be-OBut \ / \ / O O But But

Beryllium also overcomes its unsaturation by acting as a Lewis acid toward electron donor molecules. For example BeCl2 will react with solvents which have lone pairs of electron. In the reaction with ether (C2H5)2O it forms BeCl2((C2H5)2O)2. This sort of reaction is also typical of zinc halides and alkyls, and also those of magnesium and aluminium.

As stated before, certain acids in solution will react with metallic beryllium resulting in the formation of the complex [Be(H2O)4]2+ in which the water ligands are extremely tightly held. Since the Be-O interaction is so strong, there is a resultant weakening of the O-H bonds making this complex acidic by releasing H+. In aqueous solution the complex is highly dissociated and the product complex [Be(H2O)3(OH)]+ is itself unstable releasing further H+ into solution.

[Be(H2O)4]2+ → [Be(H2O)3(OH)]+ + H+

Beryllium complexes containing nitrogen ligand, such as the analogous ammonia complex [Be(NH3)4]2+, are not as stable and in aqueous solution will hydrolyse to give [Be(H2O)4]2+.

Unusual oxygen-containing beryllium complexes can be made by heating Be(OH)2 in carboxylic acids RCO2H. These form complexes with the formula Be4O(O2CR)6. They consist of an oxygen atom that is bonded to the four surrounding beryllium atoms, which are in a tetrahedral arrangement. The deprotonated carboxylic acids, or carboxylate ligands, each bond to two Be atoms forming a Be-O bond with each of their oxygen atoms. In this way, one carboxylate occupies each edge of the Be4 tetrahedron. This complex is a rare example of a tetravalent oxygen atom in a single molecule (though this readily occurs in crystal lattices of most oxides).

Compounds of the Remaining Alkaline Earth Elements

Oxides and Hydroxides

The alkaline earth metals all give oxides with the formula MO, are all white and crystalline and are obtained by heating the carbonates to very high temperature:

MCO3 → MO + CO2(g)

MgO is fairly inert but the heavier group 2 oxides will react with water to produce hydroxides. CaO is produced on a large scale for the cement industry.

MO + H2O → M(OH)2

The hydroxides are Bronsted bases. Mg(OH)2 is insoluble in water but the solubility and the strength basicity increases going down the group.


The halides of the alkaline earth metals, MX2, are easily isolated and the anhydrous salts can be obtained by heating the hydrated salts and are essentially ionic. The halides of Mg and Ca readily absorb and are soluble in water. The solubility of the halide decreases on descending the group because the hydration enthalpies decrease faster than the lattice enthalpies do. The solubilities of the fluorides of group 2 follow the opposite trend however. This is because the very small size of the fluoride ions means that in the solid state the much larger M2+ ions are more and more in contact with each other as their size increases causing a faster lowering of the lattice energies.

Oxo Salts

Oxo salts are compounds where the negative counter ion has the general formula XOnm-. The magnesium and calcium salts are often hydrates where the lattice in the solid states incorporates water molecules. The carbonates of the metals, MCO3, are all largely insoluble and the solubility decreases on going down the group. The same trend in solubility applies to the sulphates, MSO4, but magnesium sulphate is readily soluble in water. The rest of the group 2 sulphates are very insoluble, however. Other alkaline earth oxo salts are the nitrates, M(NO3)2. These can be formed the reaction of nitric acid, HNO3, with the metal hydroxide, M(OH)2.

M(OH)2 + 2 HNO3 → M(NO3)2 + 2 H2O


Only magnesium and calcium show any real ability to form complexes. They will form complexes with water, for example [Mg(H2O)6]2+, which exists in many hydrates of magnesium salts, and ammonia in [Mg(NH3)5]2+. Unlike [Be(H2O)4]2+, [Mg(H2O)6]2+ is not acidic and the water can be removed by dehydration. The halides will also form complexes with electron donating solvents. MgBr2((C2H5)2O)2 in from MgBr2 in ether. The ions also form complexes with crown ethers and cryptands in the same way as the group 1 metal ions. Ca2+ form a strong complex with the chelating ligand ethylenediaminetetraacetate (EDTA, see figure 3) which is able to interact with the ion through both nitrogen atoms and all four CO2- carboxylate groups.

Ca2+ + EDTA4- → [Ca(EDTA)]2-
Figure 3: Structure of the ligand EDTA4-. O- O- | | O=C-CH2 H2C-C=O \ /  :N-CH2-CH2-N: / \  O=C-CH2 H2C-C=O | | O- O-

Other Compounds

The metal also form compounds with most elements such as sulphur, silicon and phosphorus. These are mostly ionic and with hydrolyse with water. They will form hydrides, MH2, which are ionic, apart from MgH2 which is largely covalent as with BeH2 but isn't polymeric.

Organometallic Compounds of the Alkaline Earth Metals

As was stated earlier, beryllium organometallic compounds like the polymeric dimethyl compound, [Be(CH3)2]n (see figure 4). Here, so called three centre-two electron Be-C-Be bonds hold the chain together bridging by bridging adjacent Be atoms.

Figure 4: Structure of [Be(CH3)2]n. H H H \|/  C \ / \ / … Be Be … / \ / \   C /|\  H H H

The magnesium analogue of this compound, Mg(CH3)2, is also polymeric and can be made by the reaction of magnesium metal with dimethyl mercury:

Hg(CH3)2 + Mg → Mg(CH3)2 + Hg

The most famous and most widely studied organometallic compounds of the alkaline earth metals are the Grignard reagent. These are organomagnesium compounds derived from reacting magnesium metal with organic halides, RX, and are used in synthetic reactions in the laboratory as sources of R-. The halide, X, is usually bromide or chloride and the organic group, R, can be an alkyl or an aromatic group like a benzene ring.

Mg + R-Br → R-Mg-Br

Calcium, strontium and barium do have some organometallic compounds but these are extremely reactive and so are hard to isolate. This is because they are much more electropositive than Be and Mg and so much less likely to form covalent bonds so R- will be much more reactive. A useful compound of calcium is calcium carbide CaC2. This contains the ion C≡C2- which reacts with water and is a source of ethyne, otherwise known as 'acetylene'.

CaC2 + 2 H2O → Ca(OH)2 + HC≡CH
1Where the anionin the acid won't act as a ligand to Be2An Angstrom, Å, is a unit of length equal to 1 x 10-10 metres.

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