Reactivity of Alkaline Earth metals, Chemistry tutorial

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, M2+, 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.   

2193_Rock salt (NaC1) structure.jpg

Fig: Rock salt (NaCl) structure

32_Wurtzite (ZnS) structure.jpg

Fig: Wurtzite (ZnS) structure

583_Beryllium chloride.jpg

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.

934_Beryllium hydride polymer.jpg

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 ≡ C2- 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 )

Mg2C + 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, N3- , 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+N23-(g)

Mg(s)2+ N23- (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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