Compounds of alkaline earth metals

Magnesium sulphate, Epsom salt MgSO₄.7H₂O

Magensium sulphate occurs as kieserite MgSO₄.H₂O in Stassfurt (Germany) deposit or as Epsom salt in the mineral water of the Epsom springs in England.

Preparation

(i) From dolomite

The dolomite ore is boiled with dil. H₂SO₄.

CaCO₃.MgCO₃ + 2H₂SO₄ $\to$ CaSO₄ + MgSO₄ + 2H₂O + 2CO₂

The ppt of calcium sulphate are filtered off and the solution on concentration and cooling gives crystals of MgSO₄.7H₂O.

(ii) From Magnesite

The magensite ore is powdered and dissolved in dilute H₂SO₄. The resulting solution is concentrated and cooled when crystals of MgSO₄.7H₂O separate out.

MgCO₃ + H₂SO₄ $\to$ MgSO₄ + H₂O + CO₂

(iii) From Kieserite

The mineral Kieserite (MgSO₄.H₂O) is powdered and dissolved in water. The resulting solution upon concentration and cooling gives crystals of MgSO₄.7H₂O.

(iv) Laboratory Preparation

In the laboratory MgSO₄ is prepared by dissolving Mg metal or MgO or MgCO₃ with dilute H₂SO₄.

$Mg+H_{2}S{O}_4\to MgS{O}_4+H_2$

$MgO+H_{2}S{O}_4\to MgS{O}_4+H_{2}O$

$MgC{O}_3+H_{2}S{O}_4\to MgS{O}_4+C{O}_2+H_{2}O$

The resulting solution upon concentration and cooling gives crystals of MgSO₄.7H₂O.

Properties

It is deliquescent and readily dissolves in water. Hydrates with 12, 6 and 1 molecule of water of crystallisation are also known. All these hydrates are converted into the anhydrous salt, when heated to 200°C and on further heating they decompose to form the oxide. Magnesium sulphate gives rise to double salt with the alkali sulphate.

(i) Magnesium sulphate is a colourless efflorescent crystalline solid highly soluble in water.

(ii) Isomorphism

MgSO₄.7H₂O is isomorphous with ZnSO₄.7H₂O & FeSO₄.7H₂O compounds having same crystal structure are called isomorphous and the phenomenon is called Isomorphism.

(iii) Action of Heat

When heated it losses 6 molecules of water to give Magnesium sulphate monohydrate which becomes anhydrous when heated to 503 K and finally decomposes to MgO & SO₃ gas on strong heating.

$MgS{O}_4.7H_{2}O\underset{-6H_{2}O}{\overset{423K}{\to }}MgS{O}_4.H_{2}O\underset{-H_{2}O}{\overset{503K}{\to }}MgS{O}_4\underset{strong}{\overset{heating}{\to }}MgO+S{O}_3$

Uses

(i) MgSO₄ is used as purgative medicine.

(ii) It is used as mordant for cotton in dyeing industry.

(iii) It is used in preparation of fire proof textile and wood.

(iv) Anhydrous MgSO₄ is used as a drying agent in organic chemistry.

(v) It is used in preparation of platinised asbestors which is used as a catalyst in the contact process for the manufacture of H₂SO₄.

Illustration 1. What are isomorphous salts?

Solution: The salts having the similar crystal structure are called isomorphous salts.

Examples: MgSO₄.7H₂O, ZnSO₄.7H_­2O, FeSO₄.7H₂O

Oxides of Mg, Ca

MgO(Magnesia)

It is made by heating magnesite (MgCO₃).

MgCO₃ $\to$ MgO + CO₂

It is very slightly soluble in water imparting an alkaline reaction to the solution.

MgO + H₂O $\to$ Mg(OH)₂

Calcium oxide, Quick lime CaO

Preparation

It is prepared by heating limestone in a rotatory kiln at 1070 – 1270 K.

$CaC{O}_3\stackrel{1070-1270K}{\rightleftharpoons }CaO+C{O}_2;\Delta H=+179.9KJ$

The temperature should not be raised above 1270 K. Otherwise silica present as impurity in lime will combine with calcium oxide to form infusible calcium silicate.

$CaO+Si{O}_2\stackrel{Above1270K}{\to }CaSi{O}_3$

Properties

(i) It is a white amorphous solid with m.p. of 2870 K.

(ii) When exposed to atmosphere, it absorbs moisture and CO₂ forming slaked lime and calcium carbonate respectively.

CaO + H₂O (Moisture) $\to$ Ca(OH)₂

CaO + CO₂ $\to$ CaCO₃

(iii) On adding water, it produces a hissing sound a large amount of heat is evolved which converts water into steam. This process is called slaking of lime and the fine powder obtained is called slaked lime.

CaO + H₂O $\to$ Ca(OH)₂ ; DH = -63KJ

(iv) Action of acids and acidic oxides

It is a basic oxide and hence combines with acids and acidic oxides forming salts.

CaO + 2HCl $\to$ CaCl₂ + H₂O

CaO + SO₂ $\to$ CaSO₃

(v) Reaction with coke

When heated with coke in electric furnace at 2273 – 3273 K, it forms calcium carbide.

$CaO+3C\stackrel{2273-3273K}{\to }Ca{C}_2+CO$

(vi) Reaction with ammonium salt

On heating with ammonia salts, it liberates ammonia gas.

CaO + 2NH₄Cl $\to$ CaCl₂ + 2NH₃ + H₂O

Uses

(i) It is used as a building materials.

(ii) It is used for drying alcohols and non acidic gases.

(iii) It is used in the preparation of ammonia and soda lime (CaO + NaOH).

(iv) It is used as a basic lining in furnaces.

Illustration 2. Why does a piece of burning magnesium continue to burn in SO₂?

Solution: This is because the reaction of Mg with SO₂ is exothermic.

$2Mg+S{O}_2\to 2MgO+\frac{1}{8}{S}_8+Heat$

Hydroxides of Mg & Ca

Magnesium Hydroxide [Mg(OH)₂]

It is obtained by adding caustic soda solution to a solution of magnesium sulphate or chloride.

MgSO₄ + 2NaOH $\to$ Na₂SO₄ + Mg(OH)₂

Properties

(i) It is converted into its oxide on heating.

Mg(OH)₂ $\to$ MgO + H₂O

(ii) It dissolves in NH₄Cl solution easily.

Mg(OH)₂ + 2NH₄Cl $\to$ MgCl₂ + 2NH₄OH

Calcium hydroxide, Slaked lime [Ca(OH)₂]

Preparation

(i) From Quick lime

Calcium hydroxide is prepared on commercial scale by adding water to quick lime (Slaking of lime)

$CaO+H_{2}O\to Ca{\left(OH\right)}_2$

During the process of slaking, lumps of quick lime crumble to a fine power.

(ii) From calcium chloride

It is obtained by treating calcium chloride with caustic soda.

$CaC{l}_2+2NaOH\to Ca{\left(OH\right)}_2+2NaCl$

Properties

Physical

It is a white amorphous powder sparingly soluble in water, the solubility decreasing further with rise in temperature. An aqueous solution is known as lime water and a suspension of slaked lime in water is called milk of lime.

Chemical properties

(i) Action of heat

it loses water only at temperature above 700 K.

$Ca{\left(OH\right)}_2\stackrel{700K}{\to }CaO+H_{2}O$

(ii) Reaction with chlorine

It forms calcium hypochlorite a constituent of bleaching power.

$2Ca{\left(OH\right)}_2+2C{l}_2\to CaC{l}_2+Ca{\left(OCl\right)}_2+2H_{2}O$

(iii) Reaction with carbon dioxide

When CO_­2 is passed through lime water, it turns milky due to formation of insoluble calcium carbonate

$Ca{\left(OH\right)}_2+C{O}_2\to \underset{milkines}{CaC{O}_3}\downarrow +2H_{2}O$

If excess of Co₂ is passed CaCO₃ (ppt) dissolves to form soluble calcium bicarbonate due to which milkiness disappears.

$CaC{O}_3+C{O}_2+H_{2}O\to \underset{so\mathrm{lub}le}{Ca{\left(HC{O}_3\right)}_2}$

If this clear solution of calcium bicarbonate is heated, the solution again turns milky due to the decomposition of ca(HCO₃)₂ back to CaCO₃.

$Ca{\left(HC{O}_3\right)}_2\left(aq\right)\stackrel{\Delta }{\to }CaC{O}_3\left(s\right)+C{O}_2\left(g\right)+H_{2}O\left(\ell \right)$

(iv) Reaction with acids

Slaked lime being a strong base reacts with acids and acidic gases forming salts.

$Ca{\left(OH\right)}_2+2HCl\to CaC{l}_2+H_{2}O$

$Ca{\left(OH\right)}_2+S{O}_3\to CaS{O}_4+H_{2}O$

However, Ca(OH)₂ does not dissolve in dil. H₂SO₄ because the CaSO_­4 formed is sparingly soluble in water.

Uses

(i) Calcium hydroxide is used for absorbing acidic gases such as CO₂, NO₂, SO₂, SO₃ etc.

(ii) For preparing ammonia from ammonium salts.

(ii) For softening of hard water.

(iv) In the laboratory, as lime water for detection of CO₂.

(v) In white washing due to its disinfectant properties.

(vi) In the production of mortar which is used as a building material.

Illustration 3. Which is the weakest base among NaOH, Ca(OH)₂, KOH and Be(OH)₂

Solution: Be(OH)₂ is weakest base , because alkali metal hydroxides are more stronger base than alkaline earth metal hydroxides. Also basic character of hydroxides of alkaline earth metals increases down the group. So Be(OH)₂ is the weakest one.

Calcium Carbonate (CaCO₃)

It occurs in nature as marble, limestone, chalk, coral, calcite, etc. It is prepared as a white powder, known as precipitated chalk, by dissolving marble or limestone in hydrochloric acid and removing iron and aluminium present by precipitating with NH₃, and then adding ammonium carbonate to the solution; the precipitate is filtered, washed and dried.

CaCl₂ + (NH₄)₂CO₃ $\to$ CaCO₃ + 2NH₄Cl

Properties

It dissolves in water containing CO₂, forming Ca(HCO₃)₂ but is precipitated from solution by boiling.

CaCO₃ + H₂O + CO₂ $\rightleftharpoons$ Ca(HCO₃)₂

IIllustration 4. Thermal decomposition of a compound 'X' yields, a basic oxide ( Y ) and acidic oxide( Z ) simultaneously. The acidic oxide(Z) can be absorbed by alkaline KOH. What is X,Y,Z.

Solution: $\underset{{}^{\prime }{X}^{\prime }}{CaC{O}_3}$$\to$ $\underset{\begin{array}{c}Basic\\ (Y)\end{array}}{\mathrm{CaO}}+\underset{\begin{array}{c}Acidic\\ (Z)\end{array}}{C{O}_2}$

CO₂ + 2KOH $\to$ K₂CO₃ +H₂O

Magnesium Carbonate (MgCO₃)

It is obtained as magnesite in nature. It can be prepared as a white precipitate by adding sodium bicarbonate to a solution of a magnesium salt.

MgCl₂ + NaHCO₃ $\to$ MgCO₃ + NaCl + HCl

Properties

(i) It is very much more soluble in water.

(ii) It dissolves in water containing CO₂ due to formation of soluble bicarbonate.

MgCO₃ + H₂O + CO₂ $\to$ Mg(HCO₃)₂

Bicarbonates of Mg & Ca

Calcium bicarbonate [Ca(HCO₃)₂]

It is obtained when CaCO₃ is dissolved in water containing CO₂ but it remains in the solution form CaCO₃ + H₂O + CO₂ $\rightleftharpoons$Ca(HCO₃)₂.

Magnesium bicarbonate [Mg(HCO₃)₂]

Same as in Ca(HCO₃)₂

Illustration 5. NaHCO₃ and NaOH cannot exist together in solution- Why?

Solution: NaHCO₃ is an acid salt which must react with NaOH which is strong base. The reaction is as follows:

NaHCO₃ + NaOH $\to$Na₂CO₃ + H₂O

Halides of Mg & Ca

Calcium Chloride (CaCl₂×6H₂O)

It separates out as deliquescent crystals when a solution of lime or calcium carbonate in HCl is evaporated.

CaCO₃ + 2HCl $\to$ CaCl₂ + H₂CO₃

But it separates out from the reaction mixture as CaCl₂×6H₂O. The anhydrous salt is obtained on heating above 200°C.

Properties

It is a colourless, deliquescent salt, highly soluble in water. The anhydrous salt is an excellent drying agent.

Magnesium chloride (MgCl₂×6H₂O)

It is prepared in the laboratory by crystallizing a solution of the oxide, hydroxide or carbonate in dilute hydrochloric acid.

MgO + 2HCl $\to$ MgCl₂ + H₂O

Properties

It is colourless, crystalline salt, deliquescent in nature and exceedingly soluble in water.

Illustration 6. Complete the following reactions:

(i) $MgC{l}_2.6H_{2}O\stackrel{heat}{\to }$

(ii) $MgC{l}_2.6H_{2}O\underset{\Delta }{\overset{HCl(g)}{\to }}$

Solution: (i) $MgC{l}_2.6H_{2}O\stackrel{heat}{\to }MgO+2HCl+5H_{2}O$

(ii) $MgC{l}_2.6H_{2}O\underset{\Delta }{\overset{HCl(g)}{\to }}MgC{l}_2+6H_{2}O$

Plaster of paris, CaSO₄.1/2 H₂O or (CaSO₄)₂.H₂O

It occurs in nature as gypsum and the anhydrous salt as anhydride. It is prepared by precipitating a solution of calcium chloride or nitrate with dilute sulphuric acid.

The effect of heat on gypsum or the dihydrate presents a review of interesting changes. On heating the monoclinic gypsum is first converted into orthorhombic form without loss of water. When the temperature reaches 120°C, the hemihydrate or plaster of paris is the product. The latter losses water, becomes anhydrous above 200°C and finally above 400°C, it decomposes into calcium oxide.

2CaSO₄ $\to$ 2CaO + 2SO₂ $\uparrow$ + O₂ $\uparrow$

The following conditions are necessary

(i) The temperature should not be allowed to rise above 393 K because above this temperature the whole of water of crystallization is lost. The resulting anhydrous CaSO₄ is called dead burnt plaster because it does not have the properties of setting with water.

(ii) The gypsum should not be allowed to come in contact with carbon containing fuel otherwise some of it will be reduced to calcium sulphite.

Properties

It is a white powder. On mixing with 1/3^rd its weight of water, it forms a plastic mass which sets into a hard mass of interlocking crystals of gypsum within 5 to 15 minutes. It is due to this reason that it is called plaster. The addition of common salt accelerates the rate of setting, while a little borax or alum reduces it. The setting of plaster of paris is believed to be due to rehydration and its reconversion into gypsum.

$\underset{platerofparis}{2CaS{O}_4.\frac{1}{2}{H}_{2}O}+3H_{2}O\to \underset{gypsum}{2CaS{O}_4.2H_{2}O}$

Uses

(i) Plaster of pairs is used for producing moulds for pottery and ceramics & casts of statues & busts.

(ii) It is used in surgical bandages used for plastering broken or fractured bones.

(iii) It is also used in dentistry.

Industrial uses of lime and Limestone

Uses of lime

Calcium oxide is called lime or quick lime. It main industrial uses are

(i) It is used in steel industry to remove phosphates and silicates as slag.

(ii) It is used to make cement by mixing it with silica, alumina or clay.

(iii) It is used in making glass.

(iv) It is used in lime soda process for the conversion of Na₂CO₃ to NaOH & vice versa.

(v) It is used for softening water, for making slaked lime Ca(OH)­₂ by treatment with water and calcium carbide CaC₂.

Uses of Slaked lime [Ca(OH)₂]

(i) Slaked lime is used as a building material in form of mortar. It is prepared by mixing 3 – 4 times its weight of sand and by gradual addition of water. Its sets into a hard mass by loss of H₂O and gradual absorption of CO₂ from air.

(ii) In manufacture of bleaching powder by passing Cl₂ gas.

(iii) In making glass and in the purification of sugar and coal gas.

(iv) It is used in softening of hard water.

Uses of lime stone (CaCO₃)

(i) It is used as building material in form of marble.

(ii) In manufacture of quick lime.

(iii) It is used as a raw material for the manufacture of Na₂Co₃ in solvay – ammonia process.

(iv) Commercial limestone contains iron oxide, alumina, magnesia, silica & sulphur with a CaO content of 22 – 56% MgO content upto 21%. It is used as such as a fertilizer.

Cement

Cement is essentially a finely powdered mixture of calcium silicates and aluminates along with small quantities of gypsum which sets into a hard stone like mass when treated with water.

The chief compounds of cement are tricalcium silicate 3CaO.SiO₂, dicalcium silicate, 2CaO.SiO₂ and tricalcium aluminate 3Ca. Al₂O₃. Out of these tricalcium silicate is the most important since it has property of setting quickly and acquiring considerable strength within a few days. It usually constitute 50% of the cement.

Composition of Portland cement

Lime (CaO) - 50 – 60%

MgO - 2 – 3%

Silica (SiO₂) - 20 – 25%

Ferric oxide (Fe₂O₃) - 1 – 2%

Alumina (Al₂O₃) - 5 – 10%

Sulphur trioxide (SO₃) 1 – 2%

For a good quality cement, the ratio of alumina Al₂O₃ to silica (SiO₂) should lie between 2.5 & 4 while that of lime CaO to silica + alumina + ferric oxide should be as close to 2 as possible.

Manufacture of cement

Illustration 7. Differentiate between (i) quick-lime (ii) slaked lime and, (iii) lime water

Solution: (i) Quick lime is CaO and is obtained by heating CaCO₃ in a kiln at 1273 K.

$CaC{O}_3\stackrel{\Delta }{\to }CaO+C{O}_2$

(ii) Slaked lime is calcium hydroxide, Ca(OH)₂

$CaO+H_{2}O\to Ca{(OH)}_2$

(iii) Lime water is a solution of Ca(OH)₂ in water.