The s-Block Elements-(Part: II)-Class 11

The s-Block Elements
The s-Block Elements

Group-2 element: alkaline earth metals

This is the continuation of the previous post of Group-2 elements. In order to read the previous post, CLICK.


Group-2 of the modern periodic table includes beryllium(Be), magnesium(Mg), calcium(Ca), strontium(Sr), barium(Ba) and radium(Ra). These elements are called alkaline earth metals as their oxides are highly alkaline in nature. Radium is radioactive in nature.

Electromic configuration of Group 2 elements
Electronic configuration of Group 2 elements
Group 2 elements


Alkaline earth metals do not occur in the free state due to their high reactivity. The common minerals of these elements are as follows:

Common Minerals of Group 2 elements

Radium is obtained from the uranium mineral pitchblende.

Physical and atomic properties:

Some important physical and atomic properties of the elements of group-2 are given below:

Physical state:

Group-2 elements are normally silvery white ( due to presence of very few valency electrons with respect to the orbits available) and soft (less than alkali metals).

Atomic size or atomic radii:

The size of the alkaline earth metals increases down a group as the summation effect of the addition of new shells and increase in screening effect overcome the effect of the increase in nuclear charge. Hence Be<Mg<Ca<Sr<Ba<Ra. The atomic sizes of the alkaline earth metals are smaller than the corresponding alkali metals of the same period due to an increase in nuclear charge.

Ionic Radii:

Alkaline earth metals form M++ ion. Ionic radii increase down the group. The size of the ions is less than the atomic size of the respective element because of the removal of the outermost shell and the increase in the effective nuclear charge. Hence Be2+<Mg2+<Ca2+<Sr2+<Ba2+. Their ionic radii are smaller than those of group-1 elements(due to the removal of two electrons, the effective nuclear charge increases, pulling the outer orbit closer to the nucleus).


Alkaline earth metals possess higher cohesive energy than alkali metals. This is because cohesive energy depends upon the size and number of electrons involved in the metallic bond to form the crystalline lattice. Alkali metals possess a single valence electron per atom whereas alkaline earth metals possess two valence electrons per atom. Thus, the latter possesses a more packed structure than the former. In addition to it, the smaller size of alkaline earth metals with respect to the alkali metals also increases the cohesive energy. That’s why alkaline earth metals are denser and harder than alkali metals.

Down the group, the density of the alkaline earth metals decreases till calcium then it increases. The order is :Ca<Mg<Be<Sr<Ba. This is due to the fact, that due to gradual increase in size from Be to Ca cohesive energy decreases. Again from Ca to Ba due to the formation of different crystalline lattice, cohesive energy increases.

Melting and boiling point:

The melting and boiling point of the alkaline earth metals are higher than the alkali metals due to their small size and close-packed crystalline lattice.

Down the group, the melting and boiling point of the alkaline earth metals changes irregularly. This is due to different crystal structures.

Ionization potential:

The first ionisation potential of the alkaline earth metals is higher than that of the alkali metals due to small size and higher effective nuclear charge. The second ionisation potential of the alkaline earth metals is lower than that of the alkali metals because after removing the secondvalence electron, alkaline earth metals attain stable noble gas configuration but alkali metals losses its stable noble gas configuration.

 Na(g)    →     Na+(g)    →    Na2+(g)

Mg(g)     →    Mg+(g)    →    Mg2+(g)

The first and second ionisation potential of the alkaline earth metals decreases down the group as the combined effect of addition of new shell and screening effect of the inner electrons overweighs the effect of increased nuclear charge.

Metallic nature:

Due to low ionization potential, the alkaline earth metals easily lose their valence electrons. That’s why; alkaline earth metals are highly metallic in nature.

Due to small size and higher ionisation potential alkaline earth metals have less metallic nature than alkali metals.

Down the group, as the size of the alkaline earth metals increases, the tendency to lose electrons also increases. That’s why, down the group, the metallic nature of the alkaline earth metals increases.

Oxidation state:

The most common oxidation state of the alkaline earth metals is +2. Although their 2ndionisation potential. is quite high with respect to their 1st ionisation potential, yet they show +2 state because of the following reasons-

The M2+ possess the stable noble gas electronic configuration

In solid state, due to higher charge and low size of M2+ion large amount of energy is released during the formation of compounds with M2+ions than in the formation of compounds with M+ion. This high lattice energy compensates the high second ionisation potential of the alkaline earth metals.

In the aqueous state, the high heat of hydration value of M2+ions of the alkaline earth metals compensates their high second ionization potential.

That’s why, alkaline earth metals prefer to remain as divalent ions rather than monovalent ions.

Magnetic property and colour of divalent ions:

M2+ions possess noble gas electronic configuration with no unpaired electrons. Hence the salts of alkaline earth metals are diamagnetic in nature and their divalent ions are colourless.


Alkaline earth metals are good conductors of heat and electricity due to the presence of two loosely bound valence electrons that can move freely through the crystal lattice.

Flame colouration:

The salts of the alkaline earth metals (especially chloride) impart characteristic colour to the Bunsen flame.

Chloride salts of alkaline earth metals are volatile in nature, whenheated in bunsen flame, the outermost valence electrons jump to the higher state by absorbing energy. When the excited electrons return to the ground state, they emit extra energy in form of electromagnetic radiations which fall in visible light region thereby imparting characteristic colour to the flame.

Be and Mg atoms are small in size and their electrons are strongly held by the nucleus. Thus, large amount of energy is required to excite the electrons to higher energy levels. Such a high energy is not available in the bunsen flame. That’s why they do not impart any colour to the bunsen flame.

Nature of the compound formed:

Except for Be and Mg, the alkaline earth metals form ionic compounds. This tendency to form ionic compounds increases down the group as down the group, the metals become more electropositive.


Alkaline earth metals possess little tendency to attract electrons due to their strong electropositive nature. Thus, these elements possess a low electronegativity value.

Moreover, due to small size and higher nuclear charge, the electronegativity of the alkaline earth metals is higher than those of alkali metals.

Electron affinity:

Due to fully filled s-orbitals, alkaline earth metals possess a stable electronic configuration. Hence, they possess a low electron affinity value.

Reducing nature:

Due to low ionization potential, alkaline earth metals are strong reducing agents.

As the ionisation potential value of the alkaline earth metals are more than alkali metal, the former possess apparently less reducing power than the later.

Hydration energy:

Down the group, the value of the hydration energy of the alkaline earth metals ions decreases. This is because down the group, the size of the alkaline earth metals ions increases.

As the size of the alkaline earth metal ions are smaller than the alkali metal ions, the former possess high hydration energy value than the later. That’s why compounds of alkaline earth metals easily exist as hydrated salts like MgCl2.6H2O, CaCl2.6H2O etc.


The solubility of any salt depends upon lattice energy and hydration energy. A salt get dissolved only when the hydration energy is greater than its lattice energy.

Solubility of sulphates and carbonates decreases from Be to Ba as with the increase in size of the cation, the hydration energy decreases. BeSO4 and MgSO4 are soluble in water. This is because, due to the small size of Be2+ and Mg2+ion , they possess a high hydration energy. CaSO4is slightly soluble in water. SrSO4 and BaSO4 are almost insoluble in water.

Solubility of the hydroxides and fluorides increases down the group since lattice energy decreases more rapidly than their hydration energy.

Chemical properties:

Action of water:

Alkaline earth metals react apparently less vigorously with water with respect to alkali metals. The reactivity with water increases down the group.Ca ,Sr and Ba can react with cold water forming hydroxide and hydrogen due to greater negative standard potential (E0). Due to least negative standard potential (E0) among all alkaline earth metals, Be does not react with water or steam even at red hot conditions. Mg possesses an intermediate value of standard potential (E0). Thus Mg hardly reacts with cold water but decomposes when treated with boiling water.

M + 2H2O → M(OH)2 + H2↑   ( where M= Mg, Ca, Sr and Ba)

Action of moist air:

Alkaline earth metals form oxide(MO) with oxygen. The reactivity with oxygen increases down the group. Be and Mg are not affected by air due to the formation of a film of oxide on their surface. Ca and Sr easily get tarnished in air.

Ba burns when exposed to air. That’s why Ca, Sr and Sr are usually stored in paraffin.

When ignited, beryllium burns brilliantly to form beryllium oxide and nitride.

2Be + O2 → 2BeO

3Be + N2 → Be3N2

Magnesium burns with a dazzling light to form magnesium oxide and nitride.

2Mg + O2 → 2MgO

 3Mg + N2 → Mg3N2

As large size of cations can stabilize large size anions like peroxides. Thus, Sr and Ba can form peroxides.

Action of hydrogen:

Alkaline earth metals form ionic hydride by directly combining with hydrogen except Be.

M + H2 → MH2( where M= Mg, Ca, Sr or Ca)

Beryllium hydrides can be prepared by reacting beryllium chloride with LiAlH4.

 2BeCl2 + LiAlH4 → 2BeH2 + LiCl + AlCl3

BeH2 and MgH2 are covalent hydrides where the metal possesses four electrons in their valence shell. To complete the octet, these compounds exist as polymers in which the metal atom forms four three-centre two-electron (3c-2e) bonds or banana bonds.

All hydrides react with water to liberate hydrogen gas. CaH2, SrH2 and BaH2 form ionic hydrides.

Action of halogens:

Alkaline earth metals form halides with halogen at very high temperature.

 M + X2→MX2

Halides can also be prepared by reacting oxides, hydroxides and carbonates with


MO + 2HX→MX2 + H2O

M(OH)2 + 2HX → MX2 + 2H2O

MCO3 + 2HX → MX2 + CO2 + H2O

  1. Towards a particular halogen, the reactivity increases down the group.
  2. Towards a particular alkaline earth metal the reactivity of the halogens follows F>Cl>Br>I order.
  3. All the halides are ionic except beryllium halide.
  4. Except fluorides all halides of alkaline earth metals are water soluble. Their solubility decreases down the group.

Structure of BeCl2 :

In solid state, beryllium chloride possess a chain like polymeric structure due to incomplete octet of Be in BeCl2. All the bonds are covalent in nature. Here, Be-Cl length is 202pm, Cl-Be-Cl angle = 980, Be-Cl-Be angle= 820 and the distance between two Be atoms are 263pm.

In vapour state, BeCl2 exist dimer with bridge like structure. At about 9270C it dissociates to produce linear molecules. Cl-Be-Cl

Action with carbon:

Alkaline earth metalsform carbides when reacted with carbon. These carbides when reacts with water form organic compounds.

Be2C + 4H2O→2Be(OH)2+CH4



Be2C, CaC2, Mg2C3 are called methanide,acetylide and allylide respectively as they

produce methane ,acetylene and propyne.

Solubility in liquid ammonia:

Alkaline earth metals are soluble in liquid ammonia forming a blue metal-ammonia solution like alkali metals.

Some properties of metal-ammonia solution are:

  1. With the increase in the concentration the colour of the solution changes to bronze with metal lustre due to formation of metal ion cluster.
  2. The blue colour solution when evaporated, hexammoniatesM(NH3)6 are produced.
  3. The tendency to form hexammoniatesM(NH3)6decreases down the group.
  4. HexammoniatesM(NH3)6 are good conductor of heat and electricity.
  5. Hexammoniates M(NH3)6 decompose at high temperature

M(NH3)6→M(NH2)2 + 4NH3+H2

Reaction with acids:

Alkaline earth metals react with dilute acids to form corresponding salts with the liberation of hydrogen gas.

M + H2SO4 → MSO4 + H2

Only Be is the only group-2 element that can react with alkali to form hydrogen.Be + 2NaOH + 2H2O →Na2[Be(OH)4] + H2

Tendency to form complexes:

Alkaline earth metals can form complexes due to their smaller size and higher charge. Be and Mg possesses the maximum tendency to form complexes among the alkaline earth metals.

Down the group the tendency to form complexes decreases due to the decrease in    ion-dipole interaction with the metal ion.

BeF2 can form tetrahedral complex [BeF4]2- . Magnesium can form [Mg(NH3)6]2+. Chlorophyll is a complex of Mg.

General characteristics of the compounds of the alkaline earth metals


Except BeO all oxides of alkaline earth metals are ionic.The oxides are water soluble and their aqueous solution is alkaline in nature.Ba and Sr can form peroxides.

2Be + O2 → 2BeO

2Ca + O2 → 2CaO


Down the group the basic characteristics and solubility of the hydroxides increases. Due to high ionisation potential, small size of cations and strong force of attraction between M2+ and OH , alkaline earth metal hydroxides are less basic than alkali metals.Due to its small size and high ionisation potential, beryllium oxide and hydroxide is amphoteric in nature.

CaO + H2O→Ca(OH)2


The solubility of the carbonates decreases down the group.Thermal stability increases down the group.Carbonated decomposes to produce oxides and carbon dioxide.

Ca(HCO3)2  →  CaCO3 + CO2 +H2O


Sulphates are formed by reacting oxides hydroxides and carnonates with sulphuric acid.

MO + H2SO4 → MSO4 + H2O

Sr and Ba sulphates are insoluble.CaSO4 is sparingly soluble. Be and Mg sulphates are soluble.

Sulphates of alkaline earth metals when heated produce metal oxides and sulphur


MSO4→ MO + SO3

Thermal stability of the sulphates increases down the group. This is  due to the fact

that, with the increase in the size of the metal ion ,the polarizing power of the metal

ions decreases.


Except Be , halides of alkaline earth metals are non-volatile and ionic in nature. Towards a particular halogen, the reactivity increases down the group due to the increase in the electropositive nature of the metal. Towards a particular alkali metal the reactivity of the halogens follows:  F>Cl>Br>I order.The solubility in water follows the order :BeCl2< MgCl2< CaCl2< SrCl2< BaCl2. Except BeCl2 all other halides can form hydrates and are hygroscopic in nature. Except BeF2, all fluorides are insoluble in water as their lattice energy is higher than hydration energy. Except BeCl2and MgCl2 ,all chlorides of alkaline earth metals can impart characteristic colour to bunsen flame.


Nitrates are prepared by reaction carbonates , hydroxides and oxides of alkaline earth metals with nitric acid.




All nitrates are water soluble crystals. They decompose to oxide, nitrogen dioxide and oxygen when heated.

2Ca(NO3)2 →CaO+4NO2+O2

Be can also form basic nitrate [Be4O(NO3)6].


Oxalates of Ca, Sr and Ba are sparingly soluble in water. Beryllium oxalate is water soluble.

Abnormal behaviour of beryllium:

Beryllium is the first element of the alkaline earth metals. It shows anomalous behaviour in many of the cases like:

  • Be is harder while other alkaline earth metals are soft.
  • High melting and boiling point.
  • Does not react with water even at high temperature.
  • Form covalent compounds rather than ionic.
  • Oxides and hydroxides show amphoteric nature.
  • Beryllium carbide is covalent in nature and react with water to produce methane.
  • Due to absence of d-orbital, show a maximum coordination no of 4 rather than 6

The anomalous behaviour is due the following reasons:

  • Small size of Beatom and Be++ ion.
  • High polarising power of Be2+ ion.
  • Absence of vacant d-orbital in its valence shell.
  • Higherionisation potential and electronegativity compared to other alkaline earth metals.

Diagonal relationship of Be and Al:

Be resembles Al due to almost similar atomic and ionic size. Their similarities are:

  1. Similar electronegativity
  2. Forms covalent compounds, hydrated ions and complexes
  3. Chlorides acts as Lewis acid
  4. Bears dimeric bridged structure in vapour phase
  5. Resistant to acids due to formation of a protective oxide layer on their surface.
  6. Dissolves in strong alkali to form complexes.
  7. Oxides are insoluble, hard and possess high melting point.

Preparation, properties and uses of some important compounds



Laboratory preparation:

In laboratory it is prepared by dissolving MgO or MgCO3in dil.HCl. After completion of the reaction, the final solution is subjected to crystallisation.

MgO + 2HCl   → MgCl2 + H2O

MgCO3+2HCl→  MgCl2+H2O+CO2

From sea water:

When sea water is treated with lime, magnesium hydroxide is formed. The insoluble magnesium hydroxide is then dissolved in HCl to form a solution of magnesium chloride. The solution on concentration by evaporation and cooling, crystals of MgCl2.6H2O is obtained.

From carnallite:

Finely grinded carnallite(KCl. MgCl2.6H2O) is treated with hot water and then the solution is cooled. On cooling, KCl crystallise out leaving behind magnesium chloride in the mother liquor. The mother liquor then evaporated and cooled when crystals of MgCl2.6H2O separates out.


1.It is a colourless, deliquescent, crystal.

2.It is highly soluble in water.

3. It decomposes on heating to basic magnesium chloride. On strong heating it produce magnesium oxide.

MgCl2.6H2O → Mg(OH)Cl + 5H2O + HCl

Mg(OH)Cl → MgO + HCl

4.In the presence of ammonium chloride and ammonium hydroxide, magnesium reacts with sodium hydrogen phosphate to form white crystals of magnesium ammonium phosphate. This is the principle to identify magnesium by the gravimetric method.

MgCl2 + Na2HPO4 + NH4OH→ Mg(NH4)PO4 ↓ + Na2SO4 + H2O


  1. Used in making Sorel’s Cement (MgCl2.5MgO.xH2O),which is used in dental filing, cementing glass and porcelain, making artificial stones.
  2. For making various magnesium compounds.
  3. As a lubricant for cotton threads.
  4. Anhydrous magnesium chloride can be used to extract magnesium.



CaO is prepared by heating limestone in rotatory kiln at 1073K.

CaCO3 → CaO+CO2

During its preparation, the temperature is kept at a maximum of 1270K , else it will combine with the silica impurity forming calcium silicate.


i) It is a white amorphous solid. It melts at 2870K.

ii) When heated at 2270K by oxy-hydrogen flame it emits brilliant white light called limelight.

iii) It can readily absorb moisture and carbon dioxide forming hydroxide and carbonate.

iv) It reacts with water producing huge amount of heat and a hissing sound to form fine powder, slaked lime. The process is called slaking of lime.

v) The mixture of CaO and NaOH is known as soda-lime mixture. It is used in decarboxylation of carboxylic acid.

vi)Due to its basic nature, it reacts with acid and acidic oxides.

CaO + 2HCl → CaCl2 + H2O

CaO + SiO2 → CaSiO3

6CaO + P4O10→2Ca(PO4)2

CaO + SO2 → CaSO3

vi)When heated with coke, produce calcium carbide

vii) When heated with ammonium salts, ammonia gas is liberated.


  1. Used in making dyes, cement, calcium carbide, glass etc.
  2. In purification of sugar.
  3. As a constituent of mortar, flux in metallurgy, lining in furnaces.
  4. Used in drying alcohol, preparing ammonia, fertilizers, disinfections,germicides, soda lime mixture.
  5. Used in white washing of bulidings.
  6. Used in softening water.



From quicklime:

Commercially calcium hydroxide is prepared by spraying a limited amount of water to quicklime. It is a highly exothermic process where a tremendous amount of heat is produced along with hissing sound. Quick lime at first swells up, cracks, and then converts to a fine powder called slaked lime. This process is called slaking of lime.

From calcium chloride:

It can also be prepared by treating calcium chloride with sodium hydroxide.

CaCl2 + 2NaOH→ Ca(OH)2 + 2NaCl


1.It is a white amorphous powder, sparingly soluble in water.

2.Its solubility decreases with the rise in temperature.

3.When an excess amount of slaked lime is added to small amount of water, small amount of it dissolves while the rest remain as milky suspension. This milky suspension is called milk of lime.

4.When CO2 or SO2 is passed through slaked lime, a white ppt. of carbonate and sulphite is produced.This ppt. dissolves when excess of CO2 or SO2 is passed due to formation of soluble bicarbonate and bisulphite.

Ca(OH)2 + CO2 → CaCO3↓+ H2O

CaCO3 + CO2 + H2O → Ca(HCO3)2

Ca(OH)2 + SO2 → CaSO3↓+ H2O

CaCO3 + SO2 + H2O → Ca(HSO3)2

5.When heated at 700K, it produce CaO.

Ca(OH)2→ CaO + H2O

6.Bleaching powder is produced when dry slaked lime is reacted with chlorine at 400C.

Ca(OH)2+Cl2 →Ca(OCl)Cl+2H2O

7.In cold condition, when chlorine gas is passed through excess of lime water, calcium chloride and hypochloride are produced.

2Ca(OH)2 + 2Cl2 → CaCl2 + Ca(OCl)2 + 2H2O

8.When excess chlorine gas is passed through hot milk of lime, calcium chloride and

calcium chlorate are produced.

6Ca(OH)2 + 6Cl2 → 5CaCl2 + Ca(ClO3)2 + 6H2O


  1. Used in whitewashing, tanning industry, purifying sugar, manufacture of bleaching powder, ammonia, sodium carbonate, and caustic soda.
  2. Used in softening water and qualitative analysis of carbon dioxide and carbonate.
  3. In preparing soda lime which is used as an absorbent of gases and in the decarboxylation of sodium salts of fatty acids.
  4. It is used in whitewashing buildings and as a building material in form of mortar. Mortar is prepared by mixing slaked lime and sand in 1: 3 or 1: 4 ratio by weight. Then the mixture is treated with water to produce a thick paste called mortar. This mortar sets into a hard mass by absorbing carbon dioxide and losing water.



From slaked lime:

Calcium carbonate is prepared by passing carbon dioxide through slaked lime


From calcium chloride:

Calcium carbonate can also be prepared by reacting an aqueous solution of calcium chloride and sodium carbonate.

CaCl2 + Na2CO3 → CaCO3 + 2NaCl


1.It is white solid, insoluble in water.

2.It decompose to CaO and CO2 when strongly heated at 1072-1270K.


3.Produces chlorides and sulphates when reacted with HCl and H2SO4




  1. Used in the manufacture of marble, quick lime, cement and washing soda.
  2. Used in making antacids, toothpaste, as a flux, as a filler, making high quality paper.
  3. Used in Solvay process to manufacture sodium carbonate.
  4. It can also be used as building material.

Plaster of Paris


Plaster of paris is prepared by heating gypsum at 1200C in a large steel vessel provided with mechanical stirrers or rotary kiln.

2[CaSO4.2H2O] → 2CaSO4.H2O+3H2O

During this process two important conditions are taken care off-

1.At a temperature above 1200C gypsum loses its water of crystallisation and form anhydrous calcium sulphate called dead burnt. This dead burnt does not set to aa hard mass like cement when treated with water.

CaSO4.2H2O → CaSO4  +2H2O

2.The gypsum should be freed from any carbonaceous or any other reducing agents. The heating must be carried out without any contact of carbonaceous fuel inorder to avoid the reduction of calcium sulphate.

CaSO4 + 4C → CaS + 4CO↑


1.It is a white powder which set into a hard mass when mixed with water.

2CaSO4.H2O + 3H2O → 2[CaSO4.2H2O] →  2[CaSO4.2H2O]

The process is catalysed by sodium chloride but retarded by alum or borax.

2.When heated at 2000C, it forms anhydrous calcium sulphate, which is called dead plaster or dead burnt as it hardly sets.


  1. Used in ‘Plastering’ of broken or dislocated bones.
  2. Making moulds, chalks, casts in dentistry, making statues etc.

Portland cement

Cement is a finely powdered mixture of calcium silicate and aluminate along with small quantity of gypsum which sets into a hard mass when treated with water.

The name is derived from Portland stone, a type of building stone quarried on the Isle of Portland in Dorset, England.

The average composition of Portland cement is as follows:

Na2O + K2O2.0%

In superior quantity cement, the ratio of silica and alumina should be within 2.5 to 4.0. The ratio of calcium oxide to other oxides should be near about 2.

The composition of cement is expressed as follows:

  1. 2CaO.SiO2 ( Dicalcium silicate)
  2. 3CaO.SiO2 ( Tricalcium silicate)
  3. 3CaO.Al2O3 (Tricalcium aluminate)

Raw materials:

The raw materials required for the manufacture of cement are: Limestone, clay, gypsum, small amount of magnesia and iron oxide.

Manufacture of cement:

Portland cement is manufactured by two processes: wet process and dry process.

Wet process is preferred when the raw materials are soft, humid climatic condition and cheap availability of fuel. Dry process is preferred when the raw materials are hard.

The main raw materials (lime stone and clay) are crushed separately in a suitable machine. Then they are mixed in 3:1 ratio and grounded finely. This grinding is done either by dry or by wet process.

In the wet process, clay is washed with water to remove flint and other foreign substances. It is then mixed with definite amount of limestone and pulverised in a special mill. The pasty mass thus obtained is then uniformly homogenised to get raw slurry containing about 40%water.

In the dry process, raw materials are dried and mixed in desired proportions. The mixture is finely powdered and passed through 300 mesh sieves. After that the mixture is uniformly homogenised by compressed air to get raw meal.

The raw slurry or the raw meal is introduced into the upper end of a rotary kiln by means of screw conveyer.

In the upper part of the kiln, where the temperature is around 1000-1100K, the charge loses all its water.

In the middle part of the kiln, where the temperature is around 1100-1200K, lime stone decomposes to calcium oxide and carbon dioxide.

In the lower part of the kiln, where the temperature is around 1770-1870K, lime , alumina and silica combines together to form dicalcium silicate, tricalcium silicate, tricalcium aluminate, dicalcium aluminate and tetracalciumaluminoferrite. At this high temperature about 20-30% of mass melts and combines with solid mass to form greenish black pebbels called cement clinkers.

The hot clinkers are cooled by cold air and then mixed with 2-3% gypsum. Gypsum slows down the process of setting of cement and hence imparts greater strength to the cement. This mixture is then powdered and is known as Portland cement. Finally the cement is filled in polythene or jute bags.


Setting of cement:

When cement comes in contact with water, tricalcium silicate and tricalcium aluminate present in cement get hydrolysed and get separated as hydrated colloidal gel. This gel soon gets harden as crystalline hydrates. Some amount of calcium and aluminium hydroxides are also formed due to hydrolysis. Finally, a hard gel with calcium and aluminium hydroxide is formed. Calcium and aluminium hydroxide provide additional strength to the hard mass. The reactions involved during setting of cement are:

3CaO.SiO2 + H2O → Ca(OH)2 +2CaO.SiO2 + SiO2

2CaO.SiO2 + xH2O → 2CaO.SiO2 .xH2O

3CaO.Al2O3 + 6H2O → 3CaO.Al2O3 .6H2O

3CaO.Al2O3 .6H2O → 3Ca(OH)2 + 2Al(OH)3

4CaO.Al2O3 .Fe2O3 + 6H2O →  3CaO.Al2O3.6H2O + CaO.Fe2O3

Function of gypsum:

Gypsum removes tricalcium aluminate present in cement forming calcium sulphoaluminate. As a result, setting of cement takes place slowly and a much harder mass is set in.

3CaO.Al2O3 + 3CaSO4 + 2H2O → 3CaO.Al2O3.3CaSO4.2H2O