Alkaline Earth metals, Chemistry tutorial

INTRODUCTION

Groups one and two elements belong to the s-block of the periodic table. Their electronic configurations represent an outer shell of ns1 and ns2 for alkali and alkaline earth metals correspondingly. S-block elements are recognized to be very reactive metals and usually form ionic compounds. In this chapter we will be studying the elements of group 2 consisting of beryllium, magnesium, calcium, strontium, barium or radium. Elements Ca, Sr, Ba and Ra are termed alkaline earth metals, since their earths (earth is the old name for a mineral oxide) are alkaline in nature. (remember that the alkali metals are so termed because they form hydroxides which are strong alkalies). Beryllium is not counted as an alkaline earth metal because their oxide is not alkaline.

Alkaline earth metals:

The alkaline earth metals, like alkali metal are very reactive; thus, do not take place free in nature. All of them are found in form of their salts. Let us now learning their occurrences, extraction and utilizes. We shall as well go on to study their physical properties

Occurrence:

Beryllium, the 1st member of the group is found in small quantities in the silicate mineral, phenacite, Be2 SiO4, and beryll, 3BeO.Al2O3.6Si02.Magnesium (2.76%) and calcium (4.66%) are among the 8 most abundant elements in the earth's crust.

Magnesium (0.13%) is the 2nd most abundant metallic element next only t sodium (chloride) in sea water. It take places as magnesite, MgCO3; dolomite MgCa (CO3)2; kieserite, MgSO4 .H20 and carmallite, KMgCl3.6H20 in the earth's crust. Calcium occurs extensively as calcite and lime - stone (CaCO3) in many mountain ranges. Calcium and magnesium are extremely significant biologically as well. Calcium is found in the bones of animals and human beings. Magnesium is found in the green (chlorophyll) plants. Strontium (0.038%) and barium (0.039%) are much less abundant and occur as carbonates and sulphates. Such metals are well recognized since they occur as concentrated ores and are easy to extract. Radium is very scarce (10 -10 %) and it is a radio active element.

Extraction of Alkaline Earth Metals:

Such metals are extracted via electrolysis of their fused chlorides, although magnesium has been manufactured via the carbon reduction of its oxide.

Beryllium is attained via the electrolysis of molten beryllium chloride. Sodium chloride must be added to the melt as an electrolyte since BeCl2 is covalent and, thus is an extremely poor electrical conductor. During the electrolysis, the less active metal Be is produced at the cathode and Cl2 is evolved at the anode. Calcium is extracted from fused calcium chloride using a graphite anode and iron cathode. Strontium chloride or Barium chloride are fused for the extraction of strontium and barium correspondingly.

Uses of Alkaline Earth Metals:

Beryllium is utilized for making atomic fuel containers since it absorbs very few neutrons or does not become radioactive. Being transparent to X-rays it is utilized as a window material in X- ray apparatus. It has a number of utilizes as alloys, for example whenever mixed with Cu, Berylium amplifies the strength of Cu. 6 fold. Beryllium alloys are none sparking; thus, they are employed in making hand tools for use in the petroleum industry.

Magnesium, since of its lightness, is utilized as a construction alloy material, for instance in aircrafts. For this purpose it is alloyed with aluminium. Magnesium is as well utilized as a reducing agent in the extraction of several metals similar to titanium and uranium. It forms Gringard reagents RMgX, that are significant organic reagents. Barium strontium and calcium as free metals do not find extensive uses since they are very reactive. Calcium oxide (quicklime) is a constituent of glass, mortar and port land cement.

Physical properties:

The alkaline earth metals are rapid soft metals, but are harder than the corresponding Group one elements. This is since of their 2 valence electrons that participate in metallic bonding. They are good conductors of electricity. In pure form they are silver coloured, but on exposure to the environment, the silvery luster is lost, as of the formation of an oxide layer on the surface of the metal. Their physical properties are given in Table

Property

beryllium

Magnesium

       Mg

Calcium

      Ca

Strontium

     Sr

Barium

    Ba

Radium

    Ra

Electronic

configuration

[He]2s2

[He]3s2

[Ar]4s2

[Kr]5s2

[Xe]6s2

[Rn]7s2

Atomic weight

9.012

24.312

40.08

87.62

137.34

226.02

Ionic radius(pm)

31

65

99

113

135

 

Covalent radius

89

136

174

191

198

 

Boiling point (K)

3243

1380

1760

1607

1413

1700

Melting point(K)

1553

934

1118

1062

998

 700

Enthapy of hydration

-2455

-1900

-1565

-1415

-1275

 

 

(KJ mol-1)

 

 

 

 

 

 

Density(103*kg m3)

1.85

1.74

1.54

2.6

3.62

5.5

Electronegativity

1.5

1.2

1.0

1.0

0.9

0.9

Ionisation energy (KJmol-1)

900

738

590

549

502

509

 

1757

1450

1146

1064

965

975

Table: Properties of the group 2 metals

The atoms of the alkaline earth metals are smaller than those of the corresponding Group one elements. This is since of the amplify atomic number. Since of the resulting increase in effective nuclear charge, valence shell electrons are pulled in more firmly via the nucleus, thus reducing the size of the atom. Likewise, their ionic radii are as well smaller than those of Group one elements, as the removal of 2 orbital electrons enhances the effective nuclear charge even more. The elements are denser than Group one metals since they have 2 valence electrons per atom for bonding the atoms into a metallic lattice and consequently more mass can be packed into a smaller volume.

The density declines slightly on moving down the group from Be to Ca but amplifies considerably subsequently up to Ra. The atomic/ionic radii increase from Be to Ra due to the result of extra shells of electrons added. This outweighs the consequence of increased nuclear charge. Group two metals contain higher melting points whenever compared to group 1 metals. The reason being the +2 charge on the cations in the metallic lattice causing them to be more strongly attracted to the 'Sea of electrons or making it hard to pull them apart.

The 1st ionization energy of alkaline earth metals is more than that of corresponding alkali metals. This is since the alkaline earth metals contain higher nuclear charge and are smaller in size. The electrons are as a result more tightly held to the nucleus. The 2nd ionization energy of such elements is almost twice their 1st ionization energy. This is as once one electron has been removed, the effective nuclear charge felt via the orbital electrons is increased, so that the remaining electrons are more lightly held and therefore much more energy is required to remove the 2nd electron.

Though, their 2nd inonization energy is less than that of the corresponding alkali metals as of stability of a closed shell configuration of the univalent cations that are formed in the cases of the alkali earth metals.

The ionization energy of alkaline earth metals reduces on moving down the group. The metals of this group (beryllium is an exception) form ionic compounds. This is since the assembly of positive or negative ions into a symmetrical crystal lattice outcome in the liberation of huge amounts of energy.

Electropositive character and the reducing property (tendency to loose elections) increase on moving down the group. Because alkaline earth metals loose electrons effortlessly, they form divalent cations that contain noble gas structure without unpaired electrons. Their compounds are diamagnetic and colourless, except the anion is coloured. Ca, Sr and Ba compounds provide characteristics flame colourations which are utilized to identify them - Ca('brick red flame'), Sr ('crimson red flame') and also Ba ('apple green flame'). 

Solubility, Lattice Energy and Hydration Energy:

The solubility of alkaline earth metal compounds represents various interesting trends. The metal ions are easily hydrated, for example MgCl2.6H2O, CaC12.6H2O, BaCl2.2H2O. The hydrogen energies of such metals ions are much superior to those of alkali metal ions since of their smaller size and increased cationic charge. The lattice energies of alkaline earth metal salts are as well much higher than those of alkali metal salts.

Hydration and lattice energies decreases with increase in size of metal ions. Decreasing lattice energy favours increased solubility, even as decreasing hydration energy favours decreased solubility. If on moving down the group the hydration energy decreases more quickly than the lattice energy, the compound becomes less soluble. This take places with most of the compounds except for fluorides and hydroxides for instance, solubility of sulphates decreases from BeSO4   to BaSO-4. Due to their tiny ionic radii, Be2+ and Mg2+contain high hydration energies. Since of that BeSO4 and MgSO4 only slightly soluble in water, while SrSO4 and BaSO4 are about insoluble in water. In the case of fluorides and hydroxides the lattice energy decreases more quickly than the hydration energy. This basis a reverse trend instance, the fluorides and hydroxides increases in solubility on moving down the group

                       ΔHhyd                                                   ΔHlatt

 

M2+

MO

MCO3

MF3

MI1

Be

-2494

  -

-

-

-

Mg

-1921

-3923

-3178

-2906

-2292

Cu

-1577

-3517

-2986

-2610

-2058

Sf

-1443

-3312

-2718

-2459

-

Ba

-1305

-3120

-2614

-2367

-

Enthalpies of hydration,  ΔHhyd  alkaline of alkali earth metal ions M2+ and lattice energies ΔHlatt   of their oxides, carbonates fluorides and iodides in kg mol-1.

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