INTRODUCTION

We had studied several of the physical properties of the alkaline earth metals. We will be studying their chemical properties and also study the stability of oxy salts of alkaline earth metals.

Activity of alkaline earth metals:

Alkaline earth metals are less reactive than alkali metals when they are less electropositive than they are. Table shows various chemical reactions shown through alkaline earth metals

2M(s) + O₂(R) → 2MO(s)                   All burn if heated. Some MO2 formed.                      

M(s) + S(s) → MS(s)                        The suphides are insoluble, but hydrolyse if heated in water.

M(s) + 2H₂ O (1) → M (OH₂(s) + H₂ (g)   Be does not react even at red heat; Mg reacts with           

                                                              Steam only; others react with water at room temperature

M(s) + 2H⁺ (aq) → M2⁺ (aq) + H₂ (g)    Be only slowly; others more quickly. Not with Be. With

M(s) + H₂(g) → M²⁺2H^- (s)                  others at high temperatures only. With Mg under pressure.                 

M(s) + X₂(g) →  MX₂(s)                     X = halogen No polyhaliders are formed

3M(s) + N₂(g)  → M₃N₂(s)                   At red heat. Stability: Be > Mg >a (hydrolyse to NH 3 ) In

3M(s) + 2NH₃ (g) → M₃N2(s) + 3H₂(g)  liquid ammonia, Ca, Sr and Ba blue solution because of

                                                                     solvated electrons.                       

Be + 2H₂O + 2OH → [Be (OH)₄] ^2- + H₂(g)    Not with other alkali earth metals At high        

M(s) + 2C(s) → MC₂ (s)                                     temperatures Be forms Be₂c. Ionic compounds

Table: Reactions of the Group 2 metals

Reactivity of alkaline earth metals rises with atomic number down the group. Let us assume such reactivities at a time.

(a) All the metals release hydrogen form acids, even though beryllium reacts gradually. Beryllium liberates hydrogen whenever treated with sodium hydroxide solution. The reaction can be symbolized therefore:

Be + 2H₂O + 2OH^- → [Be (OH)₄]^2- + H₂(g)

This illustrates the anomaly in the behaviors of beryllium, in having amphoteric properties

(b) All the alkaline earth metals burn in oxygen to form oxides, MO. Oxides can as well be structured via thermal decomposition of MCO₃, M(OH)₂ , M(NO₃)₂ and MSO₄ . With the immunity of Beryllium oxide that is covalent, all another oxides are ionic in nature.

BeO has a wurtzite (hexagonal ZnS) structure. In this structure each ion has 4 nearest neighbours distributed tetrahedrally about it. Others contain sodium chloride type of structure, for example each metal ion, M^2+, is surrounded via 6 metal ions. CaO is prepared on a large scale through heating calcium carbonate in lime kilns and is utilized in the manufacture of Sodium Carbonate, Calcium carbide, glass, bleaching powder and cement.

(c) Barium peroxide, BaO₂ is formed via passing air over heated BaO, at 800K. Strontium peroxide SrO₂, is attained in a like way at high temperature and pressure. Calcium peroxide, CaO₂ is obtained as a hydrate through treating Ca (OH)₂ with hydrogen peroxide H₂O₂ and then dehydrating the product.

Magnesium peroxide, MgO₂ is attained only in the crude form through using hydrogen peroxide but no peroxide of beryllium is recognized. Such peroxides are ionic solids having (O-O) ^2- ion and can be considered as salts of very weak acids, H₂O₂. The peroxides on treatment with dilute acids form H₂O₂.

Alkaline earth metals react less readily with water than alkali metals to provide hydrogen or metal hydroxides. Beryllium doesn't react by water and steam even at red heat. Magnesium reacts with hot water and Ca, Sr and Ba reacts with cold water to form the analogous hydroxides. Beryllium hydroxide, Be (OH)₂  is amphoteric; the basic strength increases in the order Mg to Ba. Aqueous solutions of calcium and barium hydroxides are well-known as lime water and baryta water correspondingly. Ca (OH)₂ reacts with CO₂ to form first an insoluble CaCO₃ which further reacts with CO₂ to give soluble bicarbonate.

Ca (OH)₂  + CO₂  → CaCO₃ + H₂O

Insoluble water precipitate

CaCO₃ + CO₂+ H₂O → Ca (HCO₃)₂

Soluble

Calcium and barium bicarbonates are stable only in solution and decompose on removal of water to provide carbonates. This property of bicarbonates is the cause for stalactite formation (The downward growth of CaCO₃ formed on the roof of a cause through the tricking of water containing calcium compounds) and stalagmite (The upward growth from the floor of a cave; via the tricking of water containing calcium compounds) formation.

(e)  Metal halides, are attained through direct combination with halogens too as via the action of halogen on metals. Beryllium halides are covalent and other metal halides are ionic. Beryllium halides are hygroscopic and fume in air due to hydrolysis. They sublime and do not conduct electricity. Anhydrous beryllium halides are polymeric. Beryllium chloride vapours contain Becl₂ and (BeCl₂)₂ but the solid is polymerized and can be symbolized as (BeCl₂)_2.   

Fig: Rock salt (NaCl) structure

Fig: Wurtzite (ZnS) structure

Fig: Beryllium chloride (a) monomer (m) dimmer and (c) polymer

The halides are hygroscopic and form hydrates. CaCl₂ is a familiar drying agent and anhydrous MgCl₂ is significant in the electrolytic extraction of Mg.

(f) All the Group two elements except beryllium form hydrides, MH₂, through direct combination with hydrogen. Beryllium hydride can be formed through reducing beryllium chloride with lithium aluminium hydride LiAlH₄. All such hydrids are reducing agents that react with water to release hydrogen. Calcium strontium and barium hydrides are ionic and contain the hydride ion H^-. Beryllium and magnesium hydrides are covalent and polymeric (BeH₂)_n has an interesting structure. The polymeric solid have hydrogen bridges between beryllium atoms.

Fig: Beryllium hydride polymer

Each beryllium atom is bonded to 4 hydrogen atoms and each hydrogen atom forms 2 bonds as it bridges two Be atoms. Because Be has two valence electrons and H only one, it means that there are no enough electrons to form the usual type of bonds in that 2 electrons are shared between two atoms. Instead of this, center bonds are formed in which a 'banana - shaped' molecular orbital covers 3 atoms Be......H.....Be containing the two electrons. The monomeric molecule BeH₂ if formed with normal bonds would contain only 4 electrons in the outer shell of the beryllium atom and would be electron deficient. This would create the molecule extremely unstable. That is why BeH₂ exists as a cluster compound in that each and every atom shares its electrons with several neighbouring atoms and receives a share in their electrons in order to obtain a stable configuration.

(g) All the metals in the Mg - Ba series or their oxides react directly with carbon to provide the carbides (acetylides), MC₂. Such carbides are ionic in nature and contain a NaCl type of structure with M²⁺ replacing Na⁺ and C ≡ C^2- replacing Cl^-. Beryllium forms methanide, Be₂C, with carbon, and acetylide BeC₂ with acetylene.

Magnesium on heating with carbon forms Mg₂C₃, which is an allylide because with water it releases allylene (methyl/acetylene). The naming of carbides depends on the hydrocarbon they free on reaction with water. If acetylene is liberated it is termed acetylide. If methane is librated it is methanide etc:-

CaC₂ + 2H₂O → Ca (OH)₂   + C2H₂   (acetylene)

Be₂C + 4H₂O → 2Be (OH)₂   + CH₄   (methane )

Mg₂C₃   + 4H₂O → 2Mg (OH)₂   + CH-C≡CH  (allene)

Alkaline earth metals burn in nitrogen to form nitrides, M₃N₂ it needs a lot of energy to convert. The stable N₂ molecule into nitride ion, N^3- , and this is recovered from the very high lattice energies of the alkaline earth metal nitrides. Beryllium compound is rather volatile whereas others are not. They are all colourless crystalline solids which decompose on heating and react with water to release ammonia and form either the metal oxide or hydroxide, for example

2 Mg(s) + N₂(g) → Mg₃²⁺N₂^3-(g)

Mg_(s)²⁺ N₂^3- _(S) 6H₂O → 3Mg²⁺ (OH^-)_{2( s)} + 2NH₃(g)  

Thermal stability of oxy salts:

All Group two elements form oxy salts. The thermal stability of the oxy salts. The thermal stability of the oxy salts rises with the increase in electropositivity of the metal, it increases down the group. The salts of group two are thermally less stable than those of group one. The carbonates of alkaline earth metals are constant at room temperature. On heating, they decay into the analogous oxides and carbon dioxide.

MCO₃ → MO + CO₂

The temperatures at which the carbonates decompose are:

BeCO₃ < 400K; MgCO₃ 800K; CaCO₃ 1200K

SrCO₃,  1550K,  BaCO₃, 1650K.

The sulphates are more stable than the carbonates. On heating they decay into oxides and sulphur trioxide

MSO₄ → MO + SO₃  

The order of decomposition temperature of the sulphate is:

BeSO₄' 850K; Mg SO₄ ' 1150K; CaSO₄' 1400K; SrSO₄' 1650K.

Alkali metal nitrate decompose into nitrites on heating whereas alkaline earth metal nitrates decompose on heating to metal oxide, nitrogen dioxide and oxygen.

For instance  2Ca(NO₃)₂ → 2CaO + 4NO₂   + O₂

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