Anamalous Behaviours Of Beryllium

Anamalous Behaviours Of Beryllium

The properties of berrylium the first member of the alkaline earth metal, differ from the rest of the member. Its is mainly because of

(i) Its small size and high polarizing power.

(ii) Relatively high electronegativity and ionization energy as compared to other members.

(iii) Absence of vacant d – orbitals in its valence shell.

Some important points of difference between beryllium and other members (especially magnesium) are given below:

(i) Be is harder than other members of its group.

(ii) Be is lighter than Mg.

(iii) Its melting and boiling points are higher than those of Mg & other members.

(iv) Be does not react with water while Mg reacts with boiling water.

(v) BeO is amphoteric while MgO is weakly basic.

(vi) Be forms covalent compounds whereas other members form ionic compounds.

(vii) Beryllium carbide reacts with water to give methane whereas carbides of other alkaline earth metals gives acetylene gas.

(viii) Beryllium does not exhibit coordination number more than four as it has four orbitals in the valence shell. The other members of this group has coordination number 6.

Resemblance of Beryllium with Aluminium (Diagonal relationship)

The following points illustrate the anomalous behaviour of Be and its resemblance with Al.

(i) Unlike groups – 2 elements but like aluminium, beryllium forms covalent compounds.

(ii) the hydroxides of Be, [Be(OH)₂] and aluminium [Al(OH)₃] are amphoteric in nature, whereas those of other elements of group – 2 are basic in nature.

(iii) the oxides of both Be and Al i.e. BeO and Al₂O₃ are high melting insoluble solids.

(iv) BeCl₂ and AlCl₃ have bridged chloride polymeric structure.

(v) The salts of beryllium as well as aluminium are extensively hydrolysed.

(vi) Carbides of both the metal reacts with water liberating methane gas.

(vii) The oxides and hydroxides of both Be and Al are amphoteric and dissolve in sodium hydroxide as well as in hydrochloric acid.

(viii) Like Al, Be is not readily attacked by acids because of the presence of an oxide film.

Magnesium Metal

Magnesium occurs as magnesite MgCO₃, dolomite CaMg(CO₃)₂, Epsomite (MgSO_4.7H₂O) and carnalite K₂MgCl₄.6Hss₂O and langbeinite K₂Mg₂(SO₄)₃ deposits. The chloride and sulphate of magnesium occurs in sea water from which it being extracted on an increasing scale.

Extraction

(a) From magnesite or Dolomite

The ore is first calcined to form the oxide

MgCO₃ → MgO + CO₂

CaCO₃.MgCO₃ → CaO.MgO + 2CO₂

The metal is obtained from the oxide or the mixed oxides as follows:

(i) From MgO:

The oxide is mixed with carbon and heated in a current of chlorine gas.

MgO + C + Cl₂ → MgCl₂ + CO

The chloride thus obtained is subjected to electrolysis.

(ii) The mixed oxides [CaO.MgO] obtained from calcination of Dolomite [CaCO₃.MgCO₃] are redcued by ferrosilicon under reduced pressure above 1273 K.

2CaO + 2MgO + FeSi → 2Mg + Fe + Ca2SiO₄

(b) From Carnallite

The ore is dehydrated in a current of hydrogen chloride and the mixture of fused chloride is electrolysed.

(c) From Sea water

Sea water containing magnesium chloride is concentrated under the sun and is treated with calcium hydroxide Ca(OH)₂. Mg(OH)₂ is thus precipitated, filtered and heated to give the oxide.

The oxide so obtained is treated as in (a) (i) above and then electrolysed.

Electrolysis of Magnesium Chloride

Electrolysis of Magnesium Chlorine

A stream of coal gas is blown through the cell to prevent oxidation of Mg metal. Mg metal is obtained in liquid state which is further distilled to give pure magnesium.

Properties of Magnesium

Physical Properties

(i) Magnesium is a silvery white metal which soon becomes dull in air.

(ii) It is a light metal with a density of 1.74 g cm−3.

(iii) It is fairly malleable and ductile.

Chemical Properties

(i) Action of oxygen or air

Magnesium does not react with dry air but slowly gets tarnished in most air due to the formation of a thin film of the oxide, MgO. It burns in oxygen or air with a dazzling light.

(ii) Action of CO₂ and SO₂

Because of its great affinity for oxygen magnesium keeps on burning even in CO₂ or SO_2.

(iii) Action of nitrogen

On heating magnesium combines with nitrogen to form magnesium nitride.

Thus when magnesium burns in air both the oxide and the nitride are formed.

(iv) Action of halogens

Magnesium on heating with halogens readily forms the halides e.g.

(v) Action of water

Magnesium does not decompose water in cold but decomposes boiling water or steam.

(vi) Action of Acids

Dilute acids reacts with magnesium to produce dihydrogen.

However with conc. H₂SO₄, SO₂ is produced

(vii) Reaction with alkyl halide

Magnesium reacts with alkyl halides in dry ether to form covalent compound called Grignard reagent.

Uses of Magnesium

(i) The chief use of magnesium is in the preparation of alloys with aluminium, zinc, manganese and tin.

Duralium (Al = 95%, Cu = 4%. Mn = 0.5%, Mg = 0.5%)

Mangnalium (Al = 90% & Mg = 10%)

Duralium being light, tough and durable is used for the manufacture of airplanes and automobiles parts. Magnalium being light, tough and hard is used for making balance beams.

(ii) Magnesium burns with an intense lights, therefore, Mg (as power or ribbon) is used in flash bulbs for photography, fireworks and signals fibres.

(iii) Mg is used for ignition of thermite charge in aluminothermy.

(iv) A suspension of magnesium hydroxide known as milk of magnesium is used as an antacid for patients suffering from acidity.

(v) In preparation of Grignard reagent.

(vi) Being a reducing agent, magnesium is used in the extraction of boron and silicon from their respective oxides.

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 → 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 → 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.

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.

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.

Oxides of Mg, Ca

MgO(Magnesia)

It is made by heating magnesite (MgCO₃).

MgCO₃ → MgO + CO₂

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

MgO + H₂O → Mg(OH)₂

Calcium oxide, Quick lime CaO

Preparation

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

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.

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) → Ca(OH)₂

CaO + CO₂ → 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 → Ca(OH)₂ ; ΔH = −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 → CaCl₂ + H₂O

CaO + SO₂ → CaSO₃

(v) Reaction with coke

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

(vi) Reaction with ammonium salt

On heating with ammonia salts, it liberates ammonia gas.

CaO + 2NH₄Cl → 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.

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 → Na2SO₄ + Mg(OH)₂

Properties

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

Mg(OH)₂ → MgO + H₂O

(ii) It dissolves in NH4Cl solution easily.

Mg(OH)₂ + 2NH₄Cl → 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)

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.

Physical Properties

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.

(ii) Reaction with chlorine

It forms calcium hypochlorite a constituent of bleaching power.

(iii) Reaction with carbon dioxide

When CO₂ is passed through lime water, it turns milky due to formation of insoluble calcium carbonate

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

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.

(iv) Reaction with acids

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

However, Ca(OH)₂ does not dissolve in dil. H₂SO₄ because the CaSO₄ 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.

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 → CaCO₃ + 2NH₄Cl

Properties

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

CaCO₃ + H₂O + CO₂ Ca(HCO₃)₂

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 → 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 → 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 Ca(HCO₃)₂.

Magnesium bicarbonate [Mg(HCO₃)₂]

Same as in Ca(HCO₃)₂

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 → 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 → MgCl₂ + H₂O

Properties

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

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 → 2CaO + 2SO₂↑ + O₂↑

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 1/3rd 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.

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 constituents 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.