Chemical Properties

         (1) Formation of the oxides and hydroxides

         (i) The elements (except Ba and Ra) when burnt in air give oxides of ionic nature M²⁺O^2- which are crystalline in nature. The Ba and Ra though give peroxide. The affinity to form higher oxides increases from Be to Ra.

            2M + O₂ → 2MO    (M is Be, Mg or Ca )

            2M + O₂ → MO₂     (M is Ba or Sr)

         (ii) Their less reactivity than the alkali metals is evident by the fact that they are slowly oxidized on exposure to air, though the reactivity of these metals towards oxygen increases on moving down the group.

        (iii) The oxides of these metals are very stable due to high lattice energy.

         (iv) The oxides of the metal (except BeO and MgO) dissolve in water to form basic hydroxides and evolve a large amount of heat. BeO and MgO have high lattice energy and hence insoluble in water.

         (v) BeO dissolves both in acid and alkalies to give salts that is BeO possesses amphoteric nature.

                                      BeO+2NaOH→Na₂BeO₂+H₂O ;  BeO+2HCl→BeCl₂+H₂O

                                      ^{Sod. beryllate                             Beryllium chloride}

         (vi)The basic nature of oxides of alkaline earth metals increases from Be to Ra as the electropositive Character increases from Be to Ra.

         (vii)The affinity of these metal to react with water increases with increase in electropositive character that is Be to Ra.

         (viii) Reaction of Be with water is not certain, the magnesium reacts only with hot water, while the other metals react with cold water but slowly and less energetically than the alkali metals.

         (ix) The inertness of the Be and Mg towards water is because of the formation of protective, thin layer of the hydroxide on surface of the metals.

         (x) The basic nature of the hydroxides increase from Be to Ra. It is due to increase in ionic radius down the group which results in the decrease in strength of M -O bond in M -(OH)₂ to show more dissociation of the hydroxides and greater basic character.

         (xi) The solubility of hydroxides of alkaline earth metals is relatively less than their corresponding alkali metal hydroxides in addition, the solubility of hydroxides of alkaline earth metals increases from Be to Ba. Be (OH)₂ and Mg (OH)₂ are nearly insoluble, Ca (OH)₂ (generally called lime water) is sparingly soluble whereas Sr(OH)₂ and Ba (OH)₂ (generally called baryta water) are more soluble.

         The trend of solubility of these hydroxides depends on the values of lattice energy and the hydration energy of these hydroxides. Magnitude of the hydration energy remains almost same while lattice energy decreases appreciably down the group leading to more negative values for DH _solution  down the group.

            DH _solution = DH _{lattice energy} + DH _{hydration energy}

         More negative is DH _solution more is solubility of compounds.

         (xii) The basic character of the oxides and hydroxides of alkaline earth metals is lesser than their corresponding alkali metal oxides and hydroxides.

         (xiii) The aqueous solution of lime water [Ca(OH)₂] or baryta  water [Ba(OH)]₂ are used to qualitative identification and the quantative estimation of carbon dioxide, as both of them gives result to white precipitate with CO₂ because of formation of insoluble CaCO₃ or BaCO₃

                Ca(OH)₂ + CO₂ → CaCO₃ + H₂O  ;  Ba(OH)₂ + CO₂ → BaCO₃ + H₂O

                        ^{(white ppt)                                            (white ppt)}         

         SO₂ also give white ppt of CaSO₃ and BaSO₃ on passing through the lime water or baryta water. Though on passing CO₂ in the excess, white turbidity of the insoluble carbonates dissolve to provide a clear solution again due to the formation of soluble bicarbonates,

            CaCO₃ → H₂O + CO₂ → Ca(HCO₃)₂

         (2) Hydrides

         (i) Except Be, all alkaline earth metals form hydrides (MH₂) on heating directly with H₂ .    M+ H₂ → MH₂.

         (ii) The BeH₂ is prepared by action of the LiAlH₄ on BeCl₂

               2BeCl₂ + LiAlH₄ → 2BeH₂ + LiCl + AlCl₃.

         (iii) The BeH₂ and MgH₂ are covalent whereas the other hydrides are ionic.

         (iv) The ionic hydrides of Sr, Ca, Ba  liberate H₂ at the anode and the metal at cathode. 

                                                            Fusion

                                                  CaH₂     →                Ca²⁺ + 2H^-

            Anode :  2H^- → H₂ + 2e^-   Cathode :  Ca²⁺ + 2e^- → Ca

         (v) The stability of hydrides decreases from Be to Ba.

         (vi) The hydrides possesing higher reactivity for the water, dissolves readily and produce the hydrogen gas.

            CaH_2(s) + 2H₂O → Ca(OH) ₂ + 2H₂­

         (3) Carbonates and Bicarbonates

         (i) All these metal carbonates (MCO₃) are insoluble in neutral medium but soluble in the acid medium. These are precipitated by addition of the alkali metal or ammonium carbonate solution  to the solution of these metals.

            (NH₄)₂ CO₃ + CaCl₂ → 2NH₄Cl + CaCO₃

              Na₂CO₃ + BaCl₂ → 2NaCl + BaCO₃

         (ii) Alkaline earth metal carbonates are obtained as white precipitates when calculated amount of carbon dioxide is passed through the solution of the alkaline metal hydroxides.

            M(OH)_{2 (aq)} + CO_{2 (g)} → MCO_3(s) + H₂O_(l)

         And the sodium or the ammonium carbonate is added to solution of the alkaline earth metal salt such as CaCl₂.

         CaCl_{2 (aq)} + Na₂CO_{3 (aq)} →  CaCO_3(s) +2 NaCl_(aq)

         (iii) The solubility of carbonates of these kind of metals also decreases downward in group because of the decrease of hydration energy as the lattice energy remains almost unchanged as in case of sulphates.

         (vi) The carbonates of these metals decompose on heating to give oxides, temperature of the decomposition increasing from Be to Ba. Beryllium carbonate is unstable in nature.     

       (4) Halides

         (i) The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, MX₂. These halides can be prepared by the action of halogen acids (HX) on metals, hydroxides metal oxides, and carbonates.

         M + 2HX → MX₂ + H₂ ; MO + 2HX → MX₂ + H₂O

         M(OH)₂ + 2HX → MX₂ +2H₂O

         MCO₃ + 2HX → MX₂ + CO₂ + H₂O

         Beryllium chloride is however, conveniently obtained from oxide

         (ii) BeCl₂ is essentially covalent, the chlorides MgCl₂, CaCl₂ , SrCl₂ and BaCl₂ are ionic; the ionic character increases as the size  of the metal ion increases. The evidence is provided by the following facts,

         (a) Beryllium chloride is relatively low melting and volatile whereas BaCl₂ has high melting and stable.

         (b) Beryllium chloride is soluble in organic solvents.

         (iii) The halides of the members of this group are soluble in water and produce neutral solutions from which the hydrates such: MgCl₂ 6H₂O, CaCl₂.6H₂O. BaCl₂ 2H₂O can be crystallised. The affinity to form the hydrated halides decreases with the increasing size of metal ions.

         (iv) The compound BeCl₂ is readily hydrolysed with the water to form acid solution, BeCl₂ + 2H₂O → Be (OH)₂ + 2HCl.

         (v) The fluorides are relatively less soluble than the chlorides due to high lattice energies. Except the BeCl₂ and MgCl₂ chlorides of the alkaline earth metals impart characteristic colours to the flame.

         CaCl₂                                          SrCl₂                      BaCl₂ 

         Brick red colour        Crimson colour        Grassy green colour

         Structure of BeCl₂ : In the solid phase polymeric chain structure with three centre two electron bonding with Be-Cl-Be bridged structure is shown below,   

         In the phase of vapour it tends to form a chloro-bridged dimer which dissociates into linear triatomic monomer at the high temperature at nearly  1200 K.

         (5) Solubility in liquid ammonia :  Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form coloured solutions When such a solution is evaporated, hexammoniate,  M(NH₃)₆ is formed.

         (6) Nitrides

         (i) All the alkaline earth metals direct combine with N₂ provides nitrides, M₃N₂.

         (ii) The ease of formation of nitrides however decreases from Be to Ba.

         (iii) Nitrides hear are hydrolysed water to liberate

   NH₃, M₃N₂ + 6H₂O → 3M(OH)₂ + 2NH₃

         (7) Sulphates  

         (i) All these form sulphate of the type M SO₄ by the action of H₂ SO₄ on metals, their carbonates, oxides, or hydroxides.

         M + H₂SO₄ → MSO₄ + H₂ ;   MO+H₂SO₄ → MSO₄+H₂O

         MCO₃+ H₂SO₄ → MSO₄ + H₂O+CO₂

         M(OH)₂ + H₂SO₄ → MSO₄ + 2H₂O

         (ii) The solubility of sulphates in water decreases on moving down the group BeSO₄ and MgSO₄ are fairly soluble in water while BaSO₄ is completely insoluble. This is because of increases in lattice energy of sulphates down the group which predominates over hydration energy.

         (iii) Sulphate are quite stable to heat however reduced to sulphide on heating with carbon.

   MSO₄ + 2C → MS+2CO₂

         (8) Action with carbon :  Alkaline metals (except Mg Be,) when heated along with carbon to form carbides of the type MC₂ These carbides are also called acetylides as on hydrolysis they evolve acetylene.

   MC₂ + 2H₂O→M(OH) ₂ + C₂H₂

         (9) Action with sulphur and phosphorus : Alkaline earth metals directly combine with sulphur and phosphorus when heated to form sulphides of the type MS and phosphides of the type M₃P₂ respectively.

            M + S → MS  ;  3M + 2P → M₃P₂

         The sulphides on hydrolysis liberate gas H₂S while phosphides on hydrolysis evolve the phosphine.

         MS + dil. acid → H₂S  ;  M₃P₂ + dil. acid → PH₃

         The sulphides are phosphorescent and are decomposed by the water

      2MS + 2H₂O → M(OH) ₂ + M(HS)₂

         (10) Nitrates : Nitrates of these metals are soluble in the water. On heating them up they decompose into their corresponding oxides with the evolution of a mixture of nitrogen dioxide and oxygen.

         (11) Formation of complexes

         (i) Tendency to show complex ion formation depends upon high nuclear charge, smaller size, and vacant orbitals to accept electron. As alkaline metals do not have these characteristics and thus are unable to form complex ion.

         (ii) However, Be²⁺ on account of smaller size forms many complex such as (BeF₃)^1-, (BeF₄)^2-.

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