Earth alkaline metal ions

Earth alkaline metal ions

Earth alkaline metals together with the alkali metals form the so-called s-block metals. Earth alkaline metals have two electrons in their outer shell which is an s-orbital type. The chemistry of the metals is characterised by the loss of both electrons, which is a result of the relatively low ionisation energy (IE) of both electrons and the subsequent formation of the stable cation M²⁺, which has a noble gas configuration (Table 3.1).

Group 2 elements are all silvery-white metals with high reactivity, similar to alkali metals, but less soft and not as reactive. Earth alkaline metals can be mostly found in the earth’s crust in the form of their cations displayed in minerals and not as the elemental metal, as these are very reactive. For example, beryllium principally occurs as beryl (Be₃Al₂[Si₆O₁₈]), which is also known as aquamarine.

Magnesium can be found in rock structures such as magnesite (MgCO₃) and dolomite (MgCO₃⋅CaCO₃), and is the eighth most abundant element in the earth’s crust. Calcium is the fifth most abundant element and can be found in minerals such as limestone (CaCO₃) and its metamorphs such as chalk and marble.

Earth alkaline metals are harder and have a higher density than sodium and potassium and higher melting points. This is mostly due to the presence of two valence electrons and the resulting stronger metallic bond. Atomic and ionic radii increase within the group, and the ionic radii are significantly smaller than the atomic radii. Again, this is due to the existence of two valence electrons, which are located in the s orbital furthest from the nucleus. The remaining electrons are attracted even closer to the nucleus as a result of the increased effective nuclear charge. 

The IEs of the first two valence electrons are similar and relatively low compared to the energy needed to remove the third valence electron, which is part of a fully filled quantum shell. As a result, the dominant oxidation state of earth alkaline metals is +2.

 

Major uses and extraction

Beryllium is one of the lightest metals and therefore is used in high-speed aircrafts and missiles. Unfor-tunately, it is highly toxic, and CBD, a scarring of the lung tissue, is often seen in workers from within a beryllium-contaminated work environment.

Calcium is mostly found in limestone and its related forms, such as chalk, and marble and lime (CaO). Ca²⁺ ions are essential for living organisms, as is Mg²⁺. Magnesium is the only earth alkaline metal that is used on an industrial scale. It is used in ammunition (e.g. tracer bullets and incendiary bombs), as it burns with a very bright white glow. Magnesium alloyed with aluminium results in a low-density and strong material, which is used for lightweight vehicles and aeroplanes.

Magnesium is the only group 2 element that is extracted on a large scale. Its main source is seawater, and the metal is extracted by adding calcium hydroxide. Magnesium hydroxide precipitates, as it is less soluble in water compared to the calcium compound. Magnesium hydroxide is converted into magnesium chloride (MgCl₂), which can be subsequently electrolysed in a Down’s cell  in order to produce the pure magnesium metal (Figure 3.2).

Conversion 2HCl + Mg(OH)₂ → MgCl₂ + 2H₂O

Electrolysis: at the cathode: Mg²⁺ + 2e^–→ Mg

at the anode: 2Cl^–(l) Cl_2(g) + 2e^–

Redox: 2Mg²⁺ + 2Cl^–→ 2 Mg + Cl

Figure 3.2 Redox equation for the production of magnesium

Calcination ∶ Dolomite [CaMg(CO₃)₂] is converted into MgO and CaO

Reduction ∶ 2MgO + 2CaO + FeSi → 2Mg + Ca₂SiO₄ + Fe + 1450 K

Figure 3.3 Chemical equation for the production of magnesium

Alternatively, there is a second method called the ferrosilicon process or pigeon process. This involves the reduction of magnesium oxide, which is obtained from dolomite, with an iron–silicon alloy. The raw material has to be calcined first, which means the removal of water and carbon dioxide, as these would form gaseous by-products and would reverse the subsequent reduction (Figure 3.3).

 

Chemical properties

The chemical behaviour of alkaline earth metals is characterised by their strong reducing power, and there-fore they very easily form bivalent cations (M²⁺). The elements within group 2 become increasingly more electropositive on descending within the group.

The metals themselves are coloured from grey (Be, Mg) to silver (Ca, Sr, Ba) and are soft. Beryllium and magnesium are passivated and therefore kinetically inert to oxygen or water. The metal barium has to be stored under oil because of its reactivity. Metals such as calcium, strontium and barium react similar to sodium, but are slightly less reactive: All the metals except beryllium form oxides in air at room temperature once the reaction is started. The nitride compound is formed in the presence of nitrogen, and magnesium can burn in carbon dioxide, which means that magnesium fires cannot be extinguished by the use of carbon dioxide fire extinguishers.

2M+O₂ →2MO                              3M+N₂ →M₃N₂

2Mg_(s) + CO_2(g) → 2MgO_(s) + C_(s)

The oxides of alkaline earth metals have the general formula MO and are generally basic. Beryllium oxide (BeO) is formed by the ignition of beryllium metal in an oxygen atmosphere. The resulting solid is colourless and insoluble in water. Other group 2 oxides (MO) are typically formed by the thermal decomposition of the corresponding metal carbonate or hydroxide.

MCO₃ → MO + CO₂

Be(OH)₂ + 2(OH)⁻ → [Be(OH)₄]²⁻

Be(OH)₂ + H₂SO₄ → BeSO₄ + 2H₂O

Figure 3.4 Chemical equations showing the amphoteric nature of beryllium hydroxide

Peroxides are known for magnesium, calcium, strontium and barium but not for beryllium. The radius of the beryllium cation (Be²⁺) is not sufficient to accommodate the peroxide anion.

Beryllium reacts with aqueous alkali (NaOH) and forms beryllium hydroxide, which is an amphoteric hydroxide.

The term amphoteric describes compounds that can act as an acid and a base.

Beryllium hydroxide reacts with a base with the formation of the corresponding beryllium salt. Beryllium hydroxide can also be reacted with acids such as sulfuric acid and the corresponding salt, beryllium sulfate, is obtained (Figure 3.4).

Magnesium does not react with aqueous alkali (NaOH). The synthesis of magnesium hydroxide [Mg(OH)₂] is based on a metathesis reaction in which magnesium salts are reacted with sodium or potassium hydroxide.

Mg²⁺_(aq) + 2KOH_(aq) → Mg(OH)_2(s) + 2K⁺                         (3.1)

Calcium, strontium and barium oxides react exothermically with water to form the corresponding hydroxides:

CaO_(s) + H₂O_(l) → Ca(OH)_2(s)

SrO_(s) + H₂O_(l) → Sr(OH)_2(s)

BaO_(s) + H₂O_(l) → Ba(OH)_2(s)                                                    (3.2)

Magnesium and calcium hydroxides are sparingly soluble in water and the resulting aqueous solutions are mildly alkaline. In general, group 2 hydroxides, except Be(OH)₂, react as bases, and their water solubility and thermal stability increase within the group (Mg→Ba).

Earth alkaline halides (MCl₂) are normally found in their hydrated form. Anhydrous beryllium halides are covalent, whereas Mg(II), Ca(II), Sr(II), Ba(II) halides are ionic. As a result of the ionic bond, the later halides have typically high melting points and they are sparingly soluble in water. Additionally, MgCl₂, MgBr₂, MgI₂ are hygroscopic. Anhydrous calcium chloride also has a strong affinity for water and is typically used as a drying agent.

Hygroscopic compounds are substances that absorb water from the surrounding air but do not become a liquid.