UNIT 10 THE s-BLOCK ELEMENTS
Syllabus
• General introduction
• Electronic configuration
• Occurrence
• Anomalous properties of the first element of each group
• Diagonal relationship
• Trends in the variation of properties (such as ionization enthalpy, atomic and ionic radii)
• Trends in chemical reactivity with oxygen, water, hydrogen and halogens
• Uses
In the block-wise division of the periodic table, the four clearly identified blocks of elements are the s-block, p-block, the d-block and the f-block. In the build-up of atoms elements belonging to a particular block receive their outer most electrons in s- , p-, d-, or f-orbitals. As the s-orbital can accommodate only two electrons, two groups ( 1 and 2) belong to the s-block of the periodic table. Group 1 consists of elements : Lithium, sodium, potassium, rubidium, caesium and francium. They are collectively known as the alkali metals after the Arabic word al-quis meaning plant ashes since ashes of plants are particularly rich in the carbonates of sodium and potassium. The elements of Group 2 include beryllium, magnesium, calcium , strontium, barium and radium. These elements with the exception of beryllium are commonly known as the alkaline earth metals.
Among the alkali metals sodium and potassium are abundant, lithium , rubidium and caesium have much lower abundances. Francium is radioactive ; its longest-lived isotope 223Fr has a half-life of only 21 minutes. Of the alkaline earth metals calcium and magnesium rank fifth and sixth in abundance respectively in the earth’s crust. Strontium and barium have much lower abundances. Beryllium is rare and radium is rarest of all comprising only 1010 per cent of igneous rocks
A regular trend is observed in the physical and chemical properties of alkali metals with increasing atomic number. The loosely held s-electron in the outer most valence shell of these elements makes them the most electropositive metals which readily give ions, M+. The relationship between the elements of the alkaline earth metals is similar to that between the alkali metals. Calcium, strontium, barium, with two s-electrons in the valence shell , are also highly electropositive. They form M2+ ions.
The general electronic configuration of s-block elements is [noble gas]ns1 for alkali metals and [noble gas]ns2 for alkaline earth metals.
Lithium and beryllium , the first elements of Group 1 and 2 respectively exhibit some properties which are different from those of the other members of the respective group. In these anomalous properties they resemble the second element of the following group. Thus lithium shows similarities to magnesium and beryllium to aluminium in many of their properties. This type of diagonal similarity is commonly referred to as diagonal relationship in the periodic table.
The diagonal relationship is due to similarity in ionic sizes and / or charge/radius ratios of the elements.
GENERAL CHARACTERISTICS OF THE ALKALI METALS
The alkali metals show regular trends in their physical and chemical properties more clearly than any other group of elements in the periodic table. These trends are discussed below.
Atomic properties
1. Electronic configuration : All these elements have one valence electron (ns1) over and above the electronic configuration of noble gases. The outermost s-electron is very well screened from the nuclear charge in these elements.
Element Symbol Atomic
number Electronic
configuratio Abundance in earth's crust(ppm)
Lithium Li 3 [He] 2s1 65
Sodium Na 11 [Ne 3s1 28,300
Potassium K 19 [Ar] 4s1 25,900
Rubidium Rb 37 [Kr] 5s1 310
Caesium Cs 55 [Xe ] 6s1 7
Francium Fr 87 [Rn] 7s1 -
2. Atomic and ionic size : The alkali metals have the largest sizes in a particular period in the periodic table. With increase in atomic number , the atoms become larger. The monovalent ions (M+) are smaller than the parent atoms.
3. Ionisation enthalpy : The ionisation enthalpies of alkali metals are considerably low and decrease down the group from Li to Cs. This is because the effect of increasing size outweighs the increasing nuclear charge.
Physical properties
Atomic and Physical Properties of the alkali metals
All the alkali metals are silvery white , soft and light metals. Because of large size, these elements have low density which increases down the group from Li to Cs. However, potassium is lighter than sodium. The melting and boiling points of alkali metals are low indicating weak metallic bonding due to a single valence electron in them. The alkali metals and their salts impart characteristic colours to an oxidising flame. This is because the heat from the flame excites the outermost orbit electron to higher energy level. When the excited electron comes back to the ground state , there is emission of radiation in visible region as follows.
Metal Li Na K Rb Cs
Colour Crimson Yellow Violet Red violet Blue
nm 670.8 589.2 766.5 780 455.5
Alkali metals can therefore , be detected by the respective flame tests and can be determined by flame photometry or atomic absorption spectroscopy. These elements when irradiated with light , the light energy absorbed may be sufficient to make an atom lose electron. This property makes caesium and potassium useful in photoelectric cells.
CHEMICAL PROPERTIES
The alkali metals are highly reactive elements. The high chemical reactivity of these elements may be attributed to their low
ionisation energies and low heats of atomisation. The reactivity of alkali metals increases on going from Li to Cs.
1. Reactivity towards water : Although lithium has the most negative E, its reaction with water is considerably less vigorous than that of sodium which has the least negative E among the alkali metals. This behaviour of lithium is attributed to its small size and very high hydration energy. The reaction of sodium with water produces enough heat to melt it and hydrogen produced ignites in air. Other metals of the group with lower melting points reacts explosively with water. In all cases the reaction with water occurs as :
2 M + 2 H2O 2 M+ + 2 OH + H2 ( M = an alkali metal)
The hydroxides of alkali metals are strongly basic since they have low ionisation energies. The basic strength increases down the group because the ionisation energies of these metals tend to decrease and tendency to form positive ion increases.
2. Reactivity towards air : The alkali metals tarnish in air due to the formation of their oxides which in turn react with moisture to form hydroxides. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide O2 ion is stable only in presence of large cations such as K+, Na+ and Cs+.
4 Li + O2 2 Li2O (oxide)
2 Na + O2 Na2O2 (peroxide)
M + O2 2 MO2 (superoxide)
( M = K , Rb , Cs)
In all these oxides the oxidation state of the alkali metals is +1. Lithium shows exceptional behaviour in reacting directly with nitrogen of air to form nitride , Li3N. Because of their high reactivity towards air and water , they are normally kept in kerosene oil.
3. Reactivity towards oxygen : Alkali metals combine with oxygen upon heating to form different oxides depending upon their nature. Lithium forms a normal oxide (Li2O), sodium forms peroxide (Na2O2) while potassium and rest of the metals form superoxides (MO2 where M = K, Rb and Cs) upon heating in oxygen.
The stabilites of these oxides are linked with the relative sizes of the caions and anions involved and also upon the charges present on them. In general : A smaller cation can stabilise a smaller anion while a larger cation can stabilise a larger anion.
The size of the Li+ ion is smallest and has a strong positive field around it. It can combine only with a small anion such as oxide ion (O2) with a strong negative field around it. Na+ ion due to bigger size has a weaker positive field around it and , therefore , can stabilise peroxide ion (O22) which has also a weaker negative field around it. Thus , sodium forms peroxide (Na2O2). The remaining cations (K+, Rb+ and Cs+) are still bigger than in size and the magnitude of the positive field around them decreases in the same order. They can stabilise only an anion with a very weak negative field around it, i.e., superoxide ion (O2).
The valence bond structures of oxide, peroxide and superoxide ions are given below:
The superoxide ion has a three electron bond and due to the presence of unpaired electron, it is paramagnetic in nature. On the other hand, both oxide and peroxide ions have no unpaired electrons. These are therefore, diamagnetic in nature.
4.Reactivity towards dihydrogen : All alkali metals combine with hydrogen upon heating at about 673 K( Li at 1073 K) to form hydrides of the formula MH. All alkali metal hydrides are ionic solids with high melting points. They also react with proton donors such as alcohols , gaseous ammonia and alkynes.
The ionic character of the hydrides increases from Li to Cs.
5.Reactivity towards halogens : Alkali metals combine with halogens directly to form metal halides.
2 M + X2 2 MX
Lithium halides are some what covalent. It is because of the high polarization capability of lithium ion (the distortion of electron cloud of the anion by the cation is called polarization) . The Li+ ion is very small in size and has high tendency to distort electron cloud around the negative halide ion. Since anion with large size can be easily distorted , among halides , lithium iodide is the most covalent in nature.
With the exception of certain lithium halides, the halides of rest of metals are of ionic in nature(M+X). They have high melting and boiling points. The fused halides are good conductors of electricity in the fused state, these are used for the preparation of the alkali metals.
Order of reactivity of M : Li < Na < K < Rb < Cs Order of reactivity of X2 : F2 > Cl2 > Br2 > I2
6. Reducing property and E values : Alkali metals are strong reducing agents, lithium being the most and sodium the least powerful. The standard electrode potential (E)which measures the reducing power represents the overall change from M(s) to M+(aq). Thus the E value depends on the three parameters, i.e. sublimation, ionization and hydration enthalpies. With the small size of its ion lithium has the highest hydration enthalpy which accounts for its high negative Evalue and its reducing power.
7. Solutions of the alkali metals in liquid ammonia
The alkali metals dissolve in liquid ammonia, the solubility may be as high as 5 M. The solutions are coloured and metastable. The dilute solutions are blue but colour changes to bronze with increasing concentration. The blue colour is attributed to the presence of solvated electrons ; the reaction may be represented as :
M + (x + y ) NH3 [M(NH3)x]+ + [e(NH3)y]
In concentrated solutions, the ammoniated metal ions are bound by free unpaired electrons which have been described as ‘expanded metals’. The blue solutions are paramagnetic, whereas the bronze –coloured solutions are diamagnetic. The solutions are themselves unstable with respect to amide formation.
M+(am) + e(am) + NH3(ℓ) MNH2(am) + ½ H2 (g)
Na+ (am) + e (am) + NH3 (ℓ) NaNH2 (am) + ½ H2 (g)
( where ‘am’ denotes solution on ammonia)
However, under anhydrous conditions and in the absence of catalytic impurities such as transition metal ions, solutions can be stored for several days.
Uses
Lithium metal is used to make useful alloys, for example with lead to make ‘white metal’ bearings for motor engines.
GENERAL CHARACTERISTICS OF THE COMPOUNDS OF THE ALKALI METALS
All the common compounds of the alkali metals are generally ionic in nature. General characteristics of some of their compounds are given below.
OXIDES AND HYDROXIDES
On combustion in excess of air, lithium forms mainly the oxide, Li2 O(plus some peroxide Li2O2) , sodium forms the peroxide Na2 O2 (plus some Na2O) whilst potassium, rubidium and caesium forms superoxides MO2. Under appropriate conditions pure M2O, M2O2 and MO2 may be prepared. The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilization of large anions by larger cations through lattice effects. The oxides are easily hydrolysed by water to form the hydroxides according to the following reactions.
M2O + H2O 2 M+ + 2 OH
M2O2 + 2 H2O 2 M+ + 2 OH + O2
2 MO2 + 2 H2O 2 M+ + 2 OH + H2O2 + O2
The oxides and peroxides are colourless when pure, but super oxides are yellow or orange coloured. The super oxides are paramagnetic. Sodium peroxide is widely used as an oxidising agent in inorganic chemistry.
The hydroxides which are obtained by the reaction of the oxides with water are all crystalline solids. The alkali metal hydroxides are strongest of all bases and dissolves freely in water with evolution of much heat. Numerous hydrates have been prepared from aqueous solutions of the heavier alkali metal hydroxides, e.g. NaOH n H2O ( whre n = 1,2,3,4,5 and 7) but little is known about their structures.
HALIDES
The alkali metal halides MX (where M is an alkali metal and X is a halogen) are all high-melting, colourless crystalline solids which are conveniently prepared by the reaction of the appropriate oxide, hydroxide or carbonate with aqueous hydrohalic acid (HX). All these halides have high negative enthalpies of formation ; the Hf values for fluorides become less negative as we go down the group, whilst the reverse is true for fH for chlorides, bromides and iodides. For given metal fH always becomes less negative from the fluoride to iodide.
The melting ponts and boiling points always follow the trend :
Flouride > chloride > bromide > iodide
All the halides are soluble in water. The low solubility of LiF is due to its high lattice energy whereas the low solubility of CsI is due to smaller hydration energy of its two ions. In solubilities lithium salts resemble those of Mg2+ salts. Other halides of lithium are soluble in ethanol, acetone and ethyl acetate : LiCl is soluble in pyridine also.
SALTS OF OXO-ACIDS
Oxo-acids are those in which the acidic proton is on a hydroxyl group attached to the same atom e.g., carbonic acid , H2CO3 [ ( OC(OH)2] ; sulphuric acid , H2SO4 [ O2S(OH)2 ]. Since the alkali metals are the most electropositive metals they form salts with all the oxoacids. They are generally soluble in water and thermally stable. Their carbonate (M2CO3) and in most cases the bicarbonates (MHCO3) also are highly stable to heat. As the electropositive character increases down the group, the stability of the carbonates and bicarbonates increases. Lithium carbonate is not so stable to heat ; lithium being very small in size polarizes a large CO32 ion leading to the formation of more stable Li2O and CO2. Its hydrogencarbonate does not exist as a solid.
Problem
1. Arrange the following in order of the increasing covalent character : MCl, MBr, MF, MI (where M = alkali metal)
2. What is the oxidation state of K in K2O ?
3. The Eө for Cl2 / Cl is +1.36 , for I2/I is 0.53 for Ag+/Ag is +0.79 , Na+/Na is 2.71 and for Li+/Li is 3.04. Arrange the following ionic species in decreasing order of reducing strength. I, Ag , Cl, Li, Na.
ANOMALOUS PROPERTIES OF LITHIUM
An interesting feature of the representative elements is that the first member of each group differs from the rest present in the same group also called its congeners in many characteristics. Lithium is no exception. The main reason for such behaviour are :
• The small size of Li atom and Li+ ion.
• High ionisation energy and low electropositive character as compared to the rest of the members.
• High polarising power of Li+ ion resulting in covalent character of its compounds.
• Non-availability of d-electrons in its valence shell.
• Strong intermetallic bonding due its small size.
The anomalous behaviour of lithium with respect to the rest of the members is illustrated by the following characteristics.
1. Lithium is quite hard while other alkali metals are soft and can be easily cut with a knife.
2. It is not affected by air easily and does not lose its lustre while rest of the members get their surfaces tarnished on exposure to air.
3. It reacts with water only slowly at room temperature whereas other alkali metals do so readily and even catch fire.
4. Lithium hydroxide is a very weak base and is sparingly soluble in water. Hydroxides of its congeners are strong bases and are highly soluble in water.
5. Upon heating, lithium hydroxide decomposes to lithium oxide while the hydroxides of other alkali metals fail to decompose even on strong heating. However , they do sublime on heating.
6. Lithium carbonate decomposes upon heating to evolve CO2 gas while the carbonates of other alkali metals are stable to heat.
7. Lithium is the only alkali metal which combines with nitrogen directly upon heating to form Li3N.
8. Lithium nitrate upon heating decomposes to give Li2O, NO2 and O2 , while nitrates of other alkali metals form corresponding nitrites and evolve only oxygen gas.
9. Li2SO4 does not participate in the formation of double salts while the sulphates of other alkali metals are present in double salts such as potah alum : K2SO4. Al2 (SO4) 3. 24 H2O
10. Certain compounds of lithium such as Li2CO3, Li3PO4, LiF etc. are insoluble in water, while the similar compounds of other alkali metals readily dissolves in water.
11. Lithium forms monoxide (Li2O) with oxygen while sodium forms peroxide (Na2O2) and the other alkali metals give super oxides(KO2) with oxygen.
12. Upon heating with ammonia, lithium forms imide , Li2NH while the other alkali metals form amides, MNH2.
13. Lithium ion has a tendency to participate in complexes and also forms organolithium compounds such as C4H9Li. This property is not shown by by other alkali metals.
14. Lithium reacts with bromine slowly to give lithium bromide , while other alkali metals do so violently.
15. Lithium does not form acetylide when vapours of acetylene are passed through it while the other alkali metals form corresponding acetylides.
2 Na + HC CH NaC CNa + H2
Problem
04. What makes lithium to show properties uncommon to rest of the alkali metals ?
DIAGONAL RELATIONSHIP OF LITHIUM WITH MAGNESIUM
The elements of second period resemble in many respect to those elements of third period which are diagonally opposite to them. This resemblance is called diagonal relationship. As lithium (belonging to second period) is placed diagonally opposite to magnesium (belonging to the third period), the two show diagonal relationship and resemble in several properties. These similarities are mainly due to the similarity in sizes of their ions (Li+ = 76 pm , Mg2+ = 72 pm). The main points of resemblances are as follows :
• Both lithium and magnesium are harder and lighter than other elements in the respective groups.
• Lithium and magnesium react slowly with cold water. The oxides and hydroxides are much less soluble and their hydroxides decompose on heating. Both form a nitride by direct combination with nitrogen, Li3N and Mg3N2.
• The oxides , Li2O and MgO do not combine with excess oxygen to give peroxide or super oxide.
• The carbonates of lithium and magnesium decompose easily on heating to form oxide and CO2. Solid bicarbonates are not formed by lithium and magnesium.
• Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solutions as hydrates, LiCl. 2 H2O and MgCl2. 8 H2O.
• Both LiCl and MgCl2 are soluble in ethanol.
• Both lithium perchlorate and magnesium perchlorate are extremely soluble in alcohol.
Problem
02. When is a cation highly polarising ? Which alkali metal ion has the highest polarising power ?
THE METALS
The highly electropositive nature of these elements can be seen from the values of their reduction potentials. These values indicate that the reduction of their compounds to produce metals is difficult and requires electrolytic processes.
LITHIUM
Extraction : The important minerals of lithium are :
i) Lepidolite (Li, Na, K)2Al2 (SiO3)3 F(OH)
ii) Amblygonite , LiAl(PO4)F
The extraction involves the following two steps.
i) Preparation of lithium chloride from minerals.
ii) Electrolysis of lithium chloride.
(i) The ore is first heated to about 1373 K to convert it into more friable form, which is then washed with sulphuric acid at 523 K and is leached with water to give Li2SO4.H2O. Successive treatment with sodium carbonate and hydrochloric acid gives , first Li2CO3 and finally lithium chloride, LiCl. Alternatively lithium chloride can be obtained by calcining the washed ore with CaCO3 at 1273 K followed by leaching with water to give LiOH and then treatment with HCl.
(ii) Electrolysis of lithium chloride: The electrolyte consists of a fused mixture of 55% LiCl and 45% KCl at 723 K. On electrolysis the molten lithium containing about one percent potassium is collected.
Production of lithium by electrolysis
Properties and uses
Lithium is a soft silvery-white metal. It is harder than sodium but softer than lead. It is the smallest element in the group and has the highest ionisation enthalpy, melting point and heats of atomization. Lithium has the lowest density of any solid in room temperature
Lithium and its compounds have found variety of uses in recent times. The metal is used as a deoxidiser, as an alloying metal, specially with lead to give toughened bearings and with aluminium to give high strength aluminium alloys for aircraft construction. Its alloy (14% Li) with magnesium is extremely tough and corrosion resistant which is used for armour plate and aerospace components.
SODIUM
Extraction : Sodium occur in nature as:
(i) Rock salt deposits. : NaCl
(ii) Sea water
(iii) Sodium carbonate : Na2CO3
(iv) Chile salt petre NaNO3
(v) Albite (Soda feldspar) : NaAlSi3O8
(vi) Borax : (Na2B4O7. 10 H2O)
Sodium is obtained by the electrolytic reduction of fused mixture of sodium chloride(40%) and calcium chloride (60%) at 1123 K, in Downs cell.
Down’s cell
The cell consists of a steel tank lined with fire bricks. A circular graphite anode is placed in the centre of the cell. This is surrounded by a ring shaped steel cathode. The anode and cathode are separated by a steel gauze diaphragm . The steel gauze diaphragm prevents the contact of sodium (liberated at cathode) from the chlorine (liberated at anode) otherwise they will combine together to form sodium chloride. The anode is covered by a conical hood which provides the outlet for the escape of chlorine gas. The cathode is provided with a circular trough attached to a pipe connected to a reservoir to collect molten sodium produced in the process.
The melting point of NaCl is 1080 K, which can be lowered to 850 K by the addition CaCl2 to it. A fused mixture of 40% NaCl and 60% CaCl2 is taken in the cell and electric current is passed through it. The following reactions take place.
On account of electrolysis , sodium is liberated at the cathode. Due to high temperature in the cell , it is in the molten state and rises up in the pipe attached to the circular trough kept on the cathode and collects in the receiver. Chlorine is liberated at the anode and is drawn out from the hood kept at the anode. The sodium metal obtained by this method is 99.8% pure. It contains some calcium (less than 1%) which separates almost completely if the metal is allowed to cool slowly. Chlorine is obtained as a byproduct in this process.
USES
1. An alloy of sodium with mercury (sodium amalgam) is used in the preparation of a number of organic compounds.
2. Sodium vapour lamps are used for lighting.
3. Sodium is used as reagent to detect the presence of nitrogen, sulphur and halogens in organic compounds.
4. It is used to prepare a number of compounds like NaOH, KOH, NaCN , NaNH2 etc.
5. In the molten state it is used in nuclear reactors as heat transfer medium.
6. An alloy of sodium-potassium is used in high temperature thermometers.
SOME IMPORTANT COMPOUNDS OF SODIUM
Industrially important compounds of sodium include sodium peroxide, sodium hydroxide and sodium carbonate. The large scale production of these compounds and their uses are described below.
SODIUM CHLORIDE
The main source of sodium chloride is sea water which contains 2.7-2.9% of the salt. It also occurs as salt beds. The evaporation of sea water yields crude sodium chloride which contains calcium sulphate(CaSO4), sodium sulphate(Na2SO4) , calcium chloride(CaCl2), magnesium chloride(MgCl2) etc. as impurities. Calcium chloride and magnesium chloride are undesirable impurities because they are deliquescent. For purification, a saturated solution of the crude salt is prepared and the insoluble impurities are removed by filtration. The solution is then saturated with HCl gas. Due to common ion effect crystals of pure sodium chloride separates out. Chlorides of calcium and magnesium being more soluble remain in solution.
Sodium chloride is a white crystalline solid. It has a solubility of 36 g in 100 g of water at 273 K. The solubility does not increase much with increase in temperature.
Uses
i) It is an essential constituent of our food.
ii) It is used for the manufacture of sodium, caustic soda, chlorine, washing soda etc.
iii) It is used for salting out soap.
iv) It is used in freezing mixtures.
SODIUM PEROXIDE, Na2O2
Sodium peroxide is manufactured by heating sodium metal on aluminium trays in air (free from CO2)
2 Na + O2 (air) Na2O2
When pure it is colourless and faint yellow colour of the usual product arises from the presence of a small amount of Na2O. When it is exposed, it comes in contact with moist air and turns white due to the formation of NaOH and Na2CO3. Thus,
Sodium peroxide is a powerful oxidising agent and oxidises chromium (III)hydroxide to sodium chromate, Manganse(II) to sodium manganate and sulphides to sulphates. For example,
USES
i) Sodium peroxide is widely used as an oxidising agent in inorganic chemistry ; its reaction with organic compounds are dangerously violent.
ii) It is also used as a bleaching agent because of its oxidising property.
iii) Since it readily combines with carbon dioxide yielding sodium carbonate and oxygen, it may be used for the purification of air in confined spaces such as submarines.
iv) Sodium peroxide is used in the manufacture of dyes and many other chemicals such as benzoyl peroxide , sodium perborate etc.
Problem
05. Why is it that on being heated in excess of air K, Rb and Cs forms superoxides in preference to oxides and peroxides.
SODIUM HYDROXIDE (CAUSTIC SODA), NaOH
Sodium hydroxide is the most important alkali and is made commercially by the electrolysis of a saturated solution of sodium chloride. Two kinds of cells are used.
1. The mercury cathode cell (Castner – Kellner cell)
In this cell mercury flows along the bottom of the cell and is made cathode.
The Castner-Kellner cell
At the anode : 2 Cl Cl2 + 2 e
At the cathode : 2 Na+ + 2 e 2 Na
Na + Hg NaHg
The brine solution flows in the same direction and anode consists of a number of graphite blocks. The brine electrolyses and since, hydrogen has a high overvoltage at mercury cathode, sodium is preferentially discharged forming amalgam with mercury. The sodium amalgam flows out and is reacted with water to give NaOH.
2 NaHg + 2 H2O 2 NaOH + 2 Hg + H2
The mercury is re-circulated to the cell. Hydrogen and chlorine are the important by-products.
2. Diaphragm cell
In this type of cell alkali and chlorine are kept separated by the use of a diaphragm and on contact with a negative wire gauze , electrolysis begins.
A Diaphragm cell
Chlorine is liberated at graphite anode and sodium hydroxide is formed at the outside edges of the cathode.
3. Using Nafion membrane cell : The natural brine (NaCl) is electrolysed in a membrane cell in which anolyte and and the catholyte are separated by nafion membrane. Nafion is a copolymer of tetrafluoromethylene and pentafluorosulphonyl ethoxyether . The copolymer is supported by a teflon mesh.
Properties
1. Sodium hydroxide is a deliquescent solid and absorbs moisture and carbon dioxide , finally setting to a solid hydrated carbonate.
2. An aqueous solution of sodium hydroxide contains large concentration of hydoxyl ion (OH) and generally precipitates insoluble metal hydroxides from solutions containing metallic cations.
3. The hydroxides of aluminium , zinc lead and tin dissolve in excess of sodium hydroxide giving clear solution which can also be obtained when these metals are acted upon by concentrated solution of sodium hydroxide. For example, some reactions are as under.
USES
Sodium hydroxide finds many important uses.
1. In petroleum industry.
2. For the preparation of artificial silk.
3. For the preparation of pure fats and oils.
4. In textile industry.
5. In paper industry.
6. In soap industry.
7. As a reagent in laboratories.
8. For absorbing acidic gases.
SODIUM CARBONATE , Na2CO3 .10 H2O
Manufacture
SOLVAY PROCESS
Sodium carbonate is generally prepared by a process called Ammonia-Soda Process or Solvay’s Process as described below :
Principle : When carbon dioxide gas is bubbled through a brine solution saturated with ammonia, it results in the formation of sodium hydrogen carbonate.
NH3 + H2O + CO2 NH4HCO3
NaCl + NH4HCO3 NaHCO3 + NH4Cl
Sodium hydrogen carbonate so formed precipitates out because of the common ion effect caused due to the presence of excess of sodium chloride. The precipitated NaHCO3 is filtered off and then ignited to get sodium carbonate.
2 NaHCO3 Na2CO3 + CO2 + H2O
Details of the process
The various parts of the plant for the manufacture of Na2CO3 is shown in Fig.
Solvays’s process
The main functions of these units have been discussed below:
i) Ammonia absorber : A 30% solution of brine is saturated with ammonia. Various impurities like calcium and magnesium salts present in commercial sodium chloride precipitate out by carbon dioxide (which comes along with ammonia recovery plant).
2 NH3 + H2O + CO2 (NH4)2CO3
CaCl2 + (NH4)2CO3 CaCO3 + 2 NH4Cl
MgCl2 + (NH4)2CO3 MgCO3 + 2 NH4Cl
The ammoniated brine is filtered to remove the precipitated calcium or magnesium carbonate.
ii) Carbonation tower : It is a high tower fitted with perforated plates. Ammoniated brine solution is made to trickle down from the top of the tower , while carbon dioxide gas from the lime kiln is admitted from the base of the tower. Carbon dioxide rises through the small perforations and its interaction with ammoniated brine results in the formation of sodium hydrogen carbonate.
NH3 + H2O + CO2 NH4HCO3
NaCl + NH4HCO3 NaHCO3 + NH4Cl
iii) Filtration : The solution containing crystals of NaHCO3 is drawn off from the base of the carbonation tower and filtered to get NaHCO3.
iv) Calcination : The NaHCO3 obtained from the above step is heated strongly in a kiln to convert it into sodium carbonate.
2 NaHCO3 Na2CO3 + CO2 + H2O
The carbon dioxide produced here is sent to carbonating tower.
5. Ammonia Recovery Tower :The filtrate after removal of NaHCO3 contains ammonium salts such as NH4CO3 and NH4Cl. The filtrate is mixed with Ca(OH)2 and is heated with steam in ammonia recovery tower.
NH4HCO3 NH3 + H2O + CO2
2 NH4Cl + Ca(OH)2 2 NH3 + 2 H2O + CaCl2
The mixture of ammonia and carbon dioxide gases obtained is used for saturation of brine while calcium chloride is obtained as a bye product.
6. Lime kiln : The lime stone is heated at about 1300 K to obtain carbon dioxide and calcium oxide.
CaC O3 CaO + CO2
The carbon dioxide gas goes to the carbonation tower, while lime is slaked with water in a tank known as slaker to form calcium hydroxide.
The overall reaction taking place in Solvay’s process is :
2 NaCl + CaCO3 Na2CO3 + CaCl2
Sodium carbonate is readily soluble in water and the crystals obtained from the solution are the decahydrate Na2CO3 . 10 H2O. This known as washing soda.
Properties
Sodium carbonate is a crystalline solid which exists as a decahydrate, Na2CO3 10H2O . This is called washing soda. It is readily soluble in water. On heating the decahydrate loses its water of crystallization to form monohydrate. Above 373 K , the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.
Carbonate part of sodium carbonate gets hydrolysed by water to form an alkaline solution.
Uses of Sodium carbonate
i) It is used for softening hard water.
ii) A mixture of sodium carbonate and potassium carbonate is used as fusion mixture.
iii) As an important laboratory reagent both in qualitative and quantitative analysis.
iv) It is used in paper, paint and textile industries.
v) It is used for washing purposes in laundry.
vi) It is used in the manufacture of glass, soap and caustic soda.
SODIUM HYDROGEN CARBONATE
(Baking Soda) NaHCO3
Preparation: Sodium hydrogen carbonate is obtained as an intermediate product in Solvay’s process for the manufacture of sodium carbonate. It can also be prepared by passing carbon dioxide through a solution of sodium carbonate whereby sodium hydrogen carbonate being less soluble crystallises out.
Na2CO3 + CO2 + H2O 2 NaHCO3
Uses of sodium bicarbonate
i) It is used as a component of baking powder.
ii) It is used in medicine. It acts as a mild antiseptic for skin infections.
iii) It is used in fire extinguishers.
Problem
06. What happens when KO2 reacts with water ? Write the balanced chemical equation for this reaction.
07. Predict giving reason the outcome of the reaction :
LiI + KF ………
BIOLOGICAL IMPORTANCE OF SODIUM AND POTASSIUM
A typical 70 kg man contains about 90 g sodium and 170 g of potassium compared with only 5g of iron and 0.06g copper.
Sodium ions are found primarily on the outside of cells , being located in blood plasma and in the interstitial fluid which surrounds the cells. These ions participate in the transmission of nerve signals , in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells. Sodium and potassium , although so similar chemically , differ quantitatively in their ability to penetrate cell membranes , in the transport mechanisms and in their efficiency to activate enzymes. Thus, potassium ions are most abundant cations within cell fluids , where they activate many enzymes , participate in the oxidation of glucose to produce ATP and with sodium , are responsible for transmission of nerve signals.
There is very considerable variation in the concentration of sodium and potassium ions found on the opposite sides of the cell membranes. As a typical example, in blood plasma sodium is present to the extent of 143 m mol L1 whereas the potassium level is only 5 m mol L1 within the red blood cells. These concentrations change to 10 m mol L1 (Na+) and 105 m mol L1(K+). These ionic gradients demonstrate that a discriminatory mechanism, called sodium-potassium pump, operates across the cell membranes which consumes more than one-third of the ATP used by a resting animal and about 15 kg per 24 h in a resting human.
GROUP 2 ELEMENTS
ALKALINE EARTH METALS
Group 2 of the periodic table comprises the elements beryllium, magnesium, calcium , strontium, barium and radium. These (except beryllium) are known as alkaline earth metals. Their general electronic configuration may be represented as [noble gas] ns2.
Element Symbol Electronic configuration
Beryllium Be 1s22s2
Magnesium Mg 1s22s22p63s2
Calcium Ca 1s22s22p63s23p64s2
Strontium Sr 1s22s22p63s23p63d104s2
Barium Ba 1s22s22p63s23p63d104s24p64d105s25p66s2
Radium Ra 1s22s22p63s23p63d104s24p64d104f145s25p6 5d106s26p67s2
The general relatioships throughout the group are similar to those in Group 1. The first element beryllium shows ‘diagonal relationship’ to aluminium. The following TABLE gives some atomic and physical properties of the alkaline earth metals.
Atomic and Physical Properties of the Alkaline Earth Metals
ATOMIC PROPERTIES
1. Electronic configuration : A comparison with electronic configurations of the alkali metals shows that these metals with an increased nuclear charge of one unit have ns2 configuration over and above the noble gas configuration in the respective periods. The number of valence electrons in these metals is , therefore 2.
2. Atomic and ionic sizes : The atomic and ionic radii of the alkaline earth metals are smaller than those of the alkali metals in the corresponding periods. This is due to increased nuclear charge in these elements.
3. Ionization Enthalpies : The first ionization enthalpies of the alkaline earth metals are higher than those of Group1 metals. The second ionisation enthalpies of the alkaline earth metals are smaller than those of corresponding alkali metals.
4. Hydration Enthalpies : Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group.
Be2+ > Mg2+ > Ca2+ > Sr2+ > Ba2+
The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, e.g., MgCl2 and CaCl2 exist as MgCl2 6 H2O and CaCl2 6 H2O while NaCl and KCl do not form such hydrates.
Physical properties
1. Physical appearance : These metals in general are silvery white , lustrous and relatively soft, but harder than the alkali metals. Beryllium and magnesium appear to be somewhat greyish.
2. Electrical and thermal conductivities : The alkaline earth metals like those of alkali metals have high electrical and thermal conductivities which are typical characteristics of metals.
3. Flame colour : Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours to the flame respectively. In flame the electrons are excited to higher energy levels and when they drop back to the original state, energy is emitted in the form of visible light. The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame. The flame test for Ca, Sr and Ba is helpful in their detection in qualitative analysis and estimation by flame photometry.
Chemical Reactivity
1. Towards air and water : The alkaline earth metals are reactive metals though less reactive than alkali metals. Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of a surface film of oxide. Beryllium does not react with water or steam even at red heat and does not oxidise in air below 873 K. Powdered beryllium burns brilliantly on ignition to BeO and Be3N2 . Magnesium is more electropositive and burns with dazzling brilliance in air to give MgO and Mg3N2 . Calcium , strontium and barium are readily attacked by air to form the oxide and nitride. They also react with water with increasing vigour even in cold.
2. Towards halogens : The alkaline earth metals readily react with halogens at elevated temperatures to form the halides of the type , MX2.
M + X2 MX2 ( X = F , Cl , Br , I )
Thermal decomposition of (NH4)2BeF4 is the best route for the preparation of BeF2 , and BeCl2 is conveniently made from the oxide.
3. Reactivity towards hydrogen : All the elements except beryllium combine with hydrogen upon heating to form their hydrides, MH2. BeH2 , however, can be prepared by the reaction of BeCl2 with LiAlH4.
2 BeCl2 + LiAlH4 2 BeH2 + LiCl + AlCl3
4. Reactivity towards acids : The alkaline earth metals readily react with acids liberating dihydrogen.
M + 2 HCl MCl2 + H2
5. With ammonia : Like alkali metals alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions from which ammoniates [M(NH3)6]2+ can be recovered.
6. Reducing power and E values : The near constancy of the E (M2+/M) values for Group 2 metals (TABLE) are some what similar to those for Group 1 metals. Therefore these metals are electropositive and are strong reducing agents. The particularly less negative value for Be arises from large hydration energy associated with the small size of Be2+ and the relatively large value of the enthalpy of atomization of the metal. The magnesium cation is more readily reduced than the cations of the heavier members of the group because its smaller size leads to a relatively large negative value for its hydration enthalpy.
7. Solutions in liquid ammonia : Like alkali earth metals , the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.
M + ( x + y ) NH3 [ M (NH3)x) ]2+ + 2 [ e (NH3)y ]
From these solutions , the ammoniates , [M(NH3)6]2+ can be recovered.
Uses
Beryllium is used in the manufacture of alloys. Copper-beryllium alloys are used in the preparation of high strength springs. Metallic beryllium is used for making windows of X-ray tubes. Magnesium forms alloys with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-craft construction. Magnesium (powder and ribbon) is used in flash powders and bulbs , incendiary bombs and signals. A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. Magnesium carbonate is an ingredient of tooth paste. Calcium and barium metals , owing to their reactivity with oxygen and nitrogen at elevated temperatures , have often been used to remove air from vaccum tubes. Radium salts are used in radiotherapy, for example in the treatment of cancer.
General Characteristics of Compounds of the Alkaline Earth Metals
The tendency to form dipositive ions by these metals can be explained in terms of the parameters used in Born Haber’s cycle. The formation of MX or MX3 is precluded on these considerations. For example, the enthalpy of MgCl is 125 kJ/mol and the corresponding value for MgCl2 is 642 kJ/mol. Thus the enthalpy of disproportionation reaction , 2 MgCl MgCl2 + Mg is 642 2 (125) = 392 kJ/mol. This shows that any synthetic route to get MgX in the case of alkaline earth metal will result in the formation of MgX2. Hence they are uniformally dipositive in their compounds.
With increased nuclear charge and smaller size, the alkaline earth metals form compounds which are less ionic than corresponding compounds of the alkali metals. The oxides and salts of light and small sized members such as Be and Mg are more covalent than those formed by the heavier and larger sized members (Ca, Sr, Ba). A general discussion of some of their compounds are given below.
1. OXIDES AND HYDROXIDES
All of them form oxides of the formula MO ( M = an alkaline earth metal) which except for BeO have rock salt structure. The bond in BeO is essentially covalent. The enthalpies of formation of these oxides are quite high and consequently they are very stable.
Metal oxide BeO MgO CaO SrO BaO
fHө(kJ/mol) 550 590 623 590 545
All these oxides except BeO react with water to form sparingly soluble hydroxides.
MO + H2O M(OH) 2
The solubility , thermal stability and basic character of these hydroxides increase with increasing atomic number from Mg(OH)2 to Ba(OH) 2. Beryllium hydroxide , Be(OH)2 like BeO reacts with acids as well as with alkalies and is thus amphoteric.
Be(OH) 2 + 2 OH [Be(OH) 4] 2
beryllate ion
Be(OH) 2 + 2 HCl + H 2 O [Be(OH)4]Cl2
2. HALIDES
Except for beryllium halides all other halides of these metals are ionic in nature. Beryllium halides are essentially covalent and are soluble in organic solvents. The BeX2 molecules exist only in gas phase but in the condensed phases it exists as a polymer in which beryllium is 4-cordinated as shown below.
(where X is a halogen)
Of the other halides only MgBr2 and MgI2 are soluble in organic solvents. The tendency to form halide hydrates gradually decreases ( for example, MgCl2.8 H2O, CaCl2.6 H2O, SrCl2.6 H2O and BaCl2. 2 H2O) down the group. The hydrated chlorides, bromides and iodides of calcium, strontium and barium can be dehydrated by heating but those of magnesium and beryllium suffer hydrolysis.
The fluorides of alkaline earth metals are sparingly soluble in water, the solubility increases slightly with increase in cation size. Calcium fluoride (CaF2) is the only large scale source of fluorine.
4. SALTS OF OXOACIDS
(a) Carbonates : Most beryllium salts of strong oxo-acids crystallise as soluble hydrates. Beryllium carbonate is prone to hydrolysis and can be precipitated only in an atmosphere of carbon dioxide. The carbonates of magnesium and the other alkaline earth metals are sparingly soluble in water, their thermal stability increases with increasing cationic size. All these carbonates are more soluble in the presence of carbon dioxide in water owing to the formation of hydrogen carbonate ion (HCO3) but bicarbonates have not been isolated in pure state. Calcium carbonate finds use in Solvay’s process for the manufacture of sodium carbonate, in glass making and in cement manufacture.
(b) Sulphates : The sulphates of alkaline earth metals are all white solids and stable to heat. The sulphates , BeSO4 and MgSO4 are readily soluble ; the solubility decreases from CaSO4 to BaSO4. Beryllium , magnesium and calcium sulphates crystallize in hydrated form, BeSO4 4 H2O ; MgSO4 7 H2O ; CaSO4. 2H2O. The sulphates of strontium and barium crystallize with water. The greater hydration energies of Be2+ and Mg2+ ions overcome the lattice energy factor and therefore , there sulphates are soluble.
(c) Oxalates : The oxalates of Ca, Sr, Ba are sparingly soluble and solubility increases from calcium to barium. Beryllium oxalate is highly soluble and magnesium oxalate falls in line with the rest.
(d) Nitrates : The nitrates are made by dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallizes with six molecules of water. Barium nitrate crystallises as the anhydrous salt. All of them decompose on heating , giving the oxide.
2 M(NO3)2 2 MO + 4 NO2+ O2
(M = Be, Mg, Ca, Sr or Ba)
Strontium and barium nitrates are used in pyrotechnics for giving red and green flames. From all the above general characteristics of the elements of Group 2, it appears that calcium , strontium and barium form a closely related distinct group where physical and chemical properties change systematically with increasing atomic size. Thus Ca, Sr and Ba are highly electropositive, have high negative Eө values and exhibit systematic trends in solubility of their oxo-salts.
Problem
08. Name one reagent or one operation to distinguish between :
(i) BeSO4 and BaSO4
(ii) Be(OH)2 and Ba(OH)2
09. Why does the solubility of alkaline earth metal hydroxides in water increase down the group.
10. Why does the solubility of alkaline earth metal carbonates and sulphates in water decrease down the group.
ANOMALOUS BEHAVIOUR OF BERYLLIUM
The anomalous behaviour of beryllium is reflected in the following :
1. Beryllium has exceptionally small atomic and ionic sizes and thus, does not compare with other members of the group. Because of high ionization energy and small atomic size it forms compounds which are largely covalent and its salts are easily hydrolysed.
2. Beryllium does not exhibit co-ordination number more than four as in its valence shell (n=2) there are only four orbitals. The remaining members of the group have a co-ordination number of six by making use of some d-orbitals in addition to s- and p- orbitals.
3. The oxide and hydroxide of beryllium , unlike the hydroxides of other elements in this group are amphoteric in nature
DIAGONAL SIMILARITY BETWEEN BERYLLIUM AND ALUMINIUM
The ionic radius of Be2+ is estimated to be 31 pm ; the charge /radius ratio is nearly the same as that of the Al3+ ion. Hence beryllium resembles aluminium in some ways . Some of the similarities are :
(i) The two elements have same electronegativity ( Be 1.5 ; Al 1.5) and their charge/radius ratios ( 0.064 and 0.060 respectively) are very similar indicating similar field strengths.
(ii) Both metals are fairly resistant to the action of acids due to a protective film of oxide on the surface. Both metals are acted upon by strong alkalies to form soluble complexes, beryllates [Be(OH) 4] 2 and aluminates, [Al(OH) 4].
(iii) The chlorides of both beryllium and aluminium have bridged chloride structures in vapour phase.
(iv) Salts of these metals form hydrated ions, e.g. [Be(OH2)4] 2+ and [Al(OH2)6]3+ in aqueous solutions. Because of the similarity of charge/radius ratios beryllium and aluminium ions have strong tendency to form complexes. For example, beryllium forms tetrahedral complexes such as BeF42 and [Be(C2O4)2] 2 and aluminium forms octahedral complexes like AlF63 and [Al(C2O4)3] 3.
Problem
11. Why does beryllium show similarities with aluminium ?
MAGNESIUM
Occurrence
i) Magnesite MgCO3 iv) Dolomite Ca Mg(CO3)2
ii) Epsonite MgSO4.7 H2O v) Carnallite KClMgCl2 6 H2O
iii) Langbeinite K2Mg2 (SO4)3
The chloride and sulphate of magnesium occur in sea water from which it is being extracted on a large scale.
Extraction
1. From magnesite or dolomite : The ore is calcined to form the oxide.
MgCO3 MgO + CO2
CaCO3 .MgCO3 CaO.MgO + 2 CO2
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 reduced by ferrosilicon under reduced pressure above 1273 K.
2 CaO MgO + FeSi 2 Mg + Fe + Ca2SiO4
2. From carnallite : The ore is dehydrated in a current of hydrogen chloride and the mixture of fused chloride is electrolysed .
3. From sea water : Sea water containing magnesium chloride is concentrated under sun and is treated with calcium hydroxide, Ca(OH)2. Magnesium hydroxide is thus precipitated, filtered and heated to give the oxide. The oxide so obtained is mixed with carbon and heated in a current of chlorine gas.
MgO + C + Cl2 MgCl2 + CO
The chloride thus obtained is electrolysed.
Electrolysis of magnesium chloride
Magnesium chloride obtained by any of the above methods ( 1 , 2 and 3) is fused and mixed with additional mixture of sodium chloride and calcium chloride in the temperature range of 973-1023 K. The molten mixture is electrolysed . Magnesium is liberated at the cathode and chlorine is evolved at the anode.
At cathode : Mg2+ + 2 e Mg
At anode : 2 Cl Cl2 + 2 e
Electrolysis of magnesium chloride
A stream of coal gas is blown through the cell to prevent oxidation of magnesium metal. Magnesium metal is obtained in the liquid state is further distilled to get pure magnesium.
Properties and uses
Magnesium like beryllium is a greyish white metal. It appears to be unreactive to oxygen and water due to the formation of a surface film of the oxide. Magnesium amalgam liberates hydrogen from water. The metal decomposes steam and readily reacts with acids including nitric acid. Magnesium burns in air to give the oxide , MgO and the nitride Mg3N2. It also reacts with most of the non-metals . Alkyl and aryl halides react with magnesium to give Grignard reagents.
Magnesium is the lightest constructional metal in the industry having a density less than two thirds of aluminium. It is an important alloying metal. Some of the magnesium alloys contain more than 90% Mg together with 2-9% Al and 1% Zn. These alloys are used in automobile engines and in aeroplanes. Up to 5% Mg is added to most commercial aluminium alloys (e.g. duralumin, magnalium) to improve its mechanical properties, weldability and resistance to corrosion. Magnesium is also for cathodic protection of other metals, as an oxygen scavenger, as a reducing agent in the production of certain metals such as Ti, Zr and Hf and the preparation of Grignard reagent.
Alloy Constituents Application
Magnalium Al : 95%
Mg : 5% Making parts of aeroplane, balance beams and pistons of motor engines
Electron metal Mg : 95%
Zn : 4.5%
Cu : 0.5% Making parts of aeroplanes
Duralumin Al : 95%
Mg : 0.5%
Mn : 0.5%
Cu : 4% Making parts of ship
Problem
12. The second ionisation enthalpy of calcium is more than that of the first and yet calcium forms CaCl2 not CaCl ?
COMPOUNDS OF THE ALKALINE EARTH METALS
The important compounds of the alkaline earth metals include calcium oxide, calcium hydroxide, plaster of paris and magnesium sulphate. Some other substances of industrial importance are lime stone and cement.
CALCIUM OXIDE (QUICK LIME), CaO
Preparation : It is prepared on a commercial scale by heating lime stone (CaCO3) in a rotary kiln at 1070 to 1270 K.
In order to get good yield of quick lime, the following precautions must be taken.
(i) CO2 (g) formed in the reaction must be removed as soon as it is formed in order to facilitate the forward reaction.
(ii) The temperature in the kiln should not be allowed to rise beyond 1270 K otherwise silica (SiO2) present as impurity in lime stone will combine with calcium oxide.
Properties
(i) It is a white amorphous solid which melts at 2273 K.
(ii) When exposed in air , it absorbs both water vapours and carbon dioxide to form calcium hydroxide (slaked lime) and calcium carbonate.
(iii) When water is added on quick lime (CaO) the reaction is highly exothermic accompanied by hissing sound. Slaked lime is formed as the product and the reaction is called slaking of lime.
(iv) When heated with an ammonium salt, it evolves ammonia gas.
(v) It combines with solid acidic oxides at high temperatures.
CaO + SiO2 CaSiO3
6 CaO + P4O10 2 Ca3(PO4)2
USES
i) It is used as a drying agent as such or as soda lime.
ii) Large quantities of quick lime are used in the production of slaked lime (for the preparation of bleaching powder).
iii) As a constituent of mortar, it is used on a very large scale in bulding constructions.
iv) It is used as basic flux in metallurgy.
v) It is helpful in tanning industry to remove hairs from hides.
CALCIUM HYDROXIDE (SLAKED LIME) Ca(OH)2
Preparation
Calcium hydroxide is prepared by adding water to quick lime CaO. The process is known as slaking of lime.
CaO + H2O Ca(OH)2
Properties
i. Calcium hydroxide is a white amorphous powder.
ii. It is sparingly soluble in water. A suspension of slaked lime in water is called milk of lime. Lime water is a well known laboratory reagent for the detection of carbon dioxide. When carbon dioxide is passed through lime water, calcium carbonate is formed due to to which lime water turns milky.
Ca(OH) 2 + CO2 CaCO3 + H2O
On passing more carbon dioxide milkness disappears due to the
formation of calcium bicarbonate.
CaCO3 + H2O + CO2 Ca(HCO3)2
The clear solution on heating again gives milkness due to the
decomposition of calcium bicarbonate into calcium carbonate.
Ca(HCO3)2 CaCO3 + H2O + CO2
iii) Milk of lime reacts with chlorine to form hypochlorite , constituent of bleaching powder.
2 Ca(OH) 2 + 2 Cl2 Ca(ClO)2 + 2 H2O
iv) With chlorine, slaked lime gives bleaching powder.
Ca(OH) 2 + Cl2 CaOCl2 + H2O
bleaching powder
v) On heating slaked lime loses water only at temperatures > 700 K.
Uses of calcium hydroxide
Calcium hydroxide finds various uses :
1. For absorbing acid gases.
2. For preparing ammonia from ammonium chloride.
3. As lime water in laboratories.
4. It is used as a disinfectant.
5. In glass making , tanning industry, for the preparation of bleaching powder and for the purification of sugar.
6. In the production of mortar, a building material.
PLASTER OF PARIS CaSO4 ½ H2O
Hemihydrate of calcium sulphate (CaSO4)2 H2O.
It is prepared by heating gypsum (CaSO4. 2 H2O) to 390 K.
The temperature should not be allowed to rise above 390 K because above this temperature the whole of water of hydration is lost. The resulting anhydrous CaSO4 is called ‘dead burnt plaster’ because it loses the property of setting with water.
Uses
i) On mixing with water it changes into plastic mass and solidifies due to rehydration. This is called setting of plaster of paris. During the process of setting, it undergoes slight expansion (about 1%). Consequently, it produces a sharp impressions of the mould into which it is put. Therefore, it is used for producing moulds for industries such as pottery, ceramics etc.
ii) For setting broken and fractured bones in the body.
iii) For making statues, models and other decorative material.
Magnesium sulphate, Epsom salt MgSO4 7 H2O
Preparation
i) Powdered magnesite MgCO3 is dissolved in dilute sulphuric acid. The resulting solution is concentrated and cooled , when crystals of MgSO4 7 H2O appear.
The salt MgSO4 7 H2O exists as well defined transparent crystals which lose water of crystallisation on exposure to air and exhibit efflorescence. It is highly soluble in water. On heating it loses water of crystallisation leaving behind a fluffy powder.
Uses
i) It is used as a purgative in medicine.
ii) As a mordant for cotton in dyeing industries.
iii) In the manufacture paints and soaps.
iv) Anhydrous magnesium sulphate is used as a dehydrating agent.
v) In preparing fire resistant textiles.
LIME STONE CaCO3
Calcium carbonate occurs in nature in different forms such as lime stone, marble and chalk. While the first two are hard rocks, the third one is soft white amorphous mass. It also occurs as dolomite in combination with magnesium carbonate.
In the laboratory CaCO3 can be obtained by passing carbon dioxide through lime water or by adding sodium carbonate to a solution of some soluble salt of calcium.
Ca(OH)2 + CO2 CaCO3 + H2O
CaCl2 + Na2CO3 CaCO3 + 2 NaCl
The product thus obtained is known as precipitated chalk.
Lime stone on calcination gives CaO and CO2.
CaCO3 CaO + CO2
Uses
i) Lime stone is used for the preparation of lime and cement.
ii) Lime stone is also used as a flux during smelting of iron ores.
iii) Marble is used as building material.
iv) Precipitated chalk is used in medicines and in tooth pastes.
v) Marble chips are used for the preparation of carbon dioxide in the laboratory.
CEMENT
Cement is one of the most important materials used in the construction of buildings, dams , bridges , roads etc. In 1824 Joseph Aspidin , a mason working in Leeds (U.K) , for the first time , heated a mixture of lime stone, clay and water and allowed the mass to stand for some time when it hardened into a stone like mass. It resembled Portland rock which was an important naturally occuring building stone used in England during those days. He named it portland cement.
Composition of Portland cement
CaO SiO2 Al2O3 MgO Fe2O3 SO3 Na2O K2O
50-60% 5-10% 5-10% 2-3% 1-2% 1-2% 1% 1%
For a good quality of cement :
Raw materials
The main raw materials required for the manufacture are :
i) Lime stone
ii) Clay (provides both SiO2 and Al2 O3)
iii) Gypsum
Besides these , small quantities of magnesia (MgO)x and ferric oxide (Fe2O3) are also needed to impart suitable colour to the cement.
Manufacturing process
The raw materials lime stone and clay are separately crushed to fine powders. These are mixed in the ratio 3 : 1 by mass and then made into a paste with water which is known as slurry. The slurry is introduced from the upper end of cylindrical rotary kiln which is 18 to 100 metres in length and 2.5 to 4 metres in diameter as shown in Fig.
Manufacture of Portland cement
The kiln is heated by burning coal dust. The temperature inside the kiln varies from 1100 to 1800 K. In the upper part of the kiln, the temperature is low and the slurry dries up by loss of water. In the central part of the kiln, lime stone decomposes at 1300 K to form calcium oxide (lime) and CO2.
In the lower part of the kiln the temperature is of the order of 1800 K . Here, lime gets mixed with alumina and silica which are present in clay to form a mixture of silicates and aluminates of calcium as follows :
The resulting mass is very hard and has a grey look. It is called clinker. The clinker is cooled and is mixed with 2 to 3% by mass of gypsum (CaSO4 2 H2O). It is finely powdered and the mass obtained is known as portland cement. The entire process may be represented with the following flow sheet diagram.
Setting of Cement
When cement is to be used , it is mixed with water to form a gelatinous mass which slowly become very hard in which SiOSi and SiOAl chains are formed. This is called setting of cement. The constituents present in cement take different times for setting. For example, tricalcium aluminate immediately sets into hard mass and is responsible for the internal strength of the cement. The tetracalcium aluminoferate also sets but not as quickly as tricalcium aluminate. Out of the silicates, tricalcium silicate also sets quickly but acquires proper strength after few days. Dicalcium silicate takes comparatively more time for proper setting. Gypsum is added to cement in order to regulate the time of setting.
Fly ash , a waste product from steel industry has properties similar to cement. It can be added to cement to reduce the cost with out affecting its quality.
Uses
Cement has become a commodity of national necessity for any country next to iron and steel. It is used in concrete and reinforced concrete , in plastering and in the construction of bridges, dams and buildings.
BIOLOGIGAL IMPORTANCE OF MAGNESIUM AND CALCIUM
An adult body contains about 25 g of magnesium and 1200 g of calcium compared with only 5 g of iron and 0.06 g of copper. The daily requirement in the human body has been estimated to be 200-300 mg.
All enzymes that utilize ATP in phosphate transfer require magnesium as cofactor. The main pigment for the absorption of light in plants is chlorophyll which contains magnesium. About 99% of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function , interneuronal transmission, cell membrane integrity and blood coagulation. The calcium concentration in plasma is regulated at about 100 mgL1. It is maintained by two hormones ; calcitonin and parathyroid hormone. The bone is not an inert and unchanging substance, but it is continuously being solubilised and redeposited to the extent of 400 mg per day in man. All this calcium passes through the plasma.
QUESTIONS
1. What are the common physical and chemical features of alkali metals ?
2. Discuss the general characteristics and gradation in properties of alkaline earth metals.
3. Why are alkali metals not found in nature ?
4. Find the oxidation state of sodium in Na2O2.
5. Explain why sodium is less reactive than potassium.
6. Compare the alkali metals and alkaline earth metals with sodium with respect to (i) ionization energy (ii) basicity of oxides (iii) solubility of hydroxides.
7. In what way lithium shows similarity to magnesium in its chemical behaviour ?
8. Explain why alkali and alkaline earth metals cannot by chemical reduction.
9. Why potassium and caesium , rather than lithium used in photoelectric cells ?
10. When an alkali metal is dissolved in liquid ammonia , the solution can acquire different colours. Explain the reason for this type of colour change.
11. Beryllium and magnesium does not give colour to the flame while other alkaline earth metals do so. Explain.
12. Discuss the various reactions which occur in Solvay Process.
13. Potassium carbonate cannot be prepared by Solvay process.
14. Lithium carbonate decomposes on heating while sodium carbonate does not.
15. Name the alkali metals which form super oxides on heating in excess of air.
16. Starting from sodium chloride, how will you prceed to prepare (i) sodium metal (ii) sodium hydroxide (iii) sodium peroxide (iv) sodium carbonate.
17. What happens when :
(i) magnesium is burnt in air
(ii) quick lime is heated with silica.
(iii) chlorine reacts with slaked lime
(iv) calcium nitrate is heated
18. Describe the important uses of each of the following (i) caustic soda (ii) sodium carbonate (iii) quick lime
19. Draw the structures of (i) BeCl2 (in solid state) (ii) BeCl2 (in vapour state)
20. The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding compounds of magnesium and calcium are sparingly soluble. Explain.
21. Describe the importance of the following :
(i) lime stone (ii) cement (iii) Plaster of paris
22. Why are lithium compounds are commonly hydrated while those of other alkali ions usually anhydrous ?
23. Why is LiF almost insoluble in water while LiCl is soluble not only in water but also in acetone.
24. Explain the significance of sodium, potassium , magnesium and calcium as biological fluids.
25. What happens when :
(i) sodium metal dropped in water
(ii) sodium metal is heated in free supply of air
(iii) sodium peroxide dissolves in water
26. (a) Why LiF is least soluble in water among the fluorides of
alkali metal ?
(b) Justify the given order of mobilities of the alkali metal
cations in aqueous solution :
Li+ < Na+ < K+ < Rb+ < Cs+
(c) Lithium is the only alkali metal which forms a nitride directly.
27. State as to why :
(a) A solution of Na2CO3 is alkaline
(b) Alkali metals are prepared by the electrolysis of their fused
chlorides.
(c) Sodium is found to be more useful than potassium
28. Write the balanced equation for the reactions between :
(i) Na2O2 and water
(ii) KO2 and water
(iii) Na2O and CO2
29. How would you explain :
(i) BeO is insoluble in water , while BeSO4 is soluble
(ii) Ba(OH)2 is soluble in water while BaSO4 is almost
insoluble
(iii) LiI is more soluble than KI in ethanol
30. Which of the following alkali metals is having least melting point ?
(i) Na (ii) K (iii) Rb (iv) Cs
31. Which of the following give hydrated salts ?
(i) Li (ii) Na (iii) K (iv) Cs
32. Thermally the most stable alkaline earth metal carbonate is :
(a) MgCO3 (b) CaCO3 (c) SrCO3 (d) BaCO3
33. List some important ores of sodium and potassium.
34. How is sodium isolated by Down’s process ? Sketch the diagram for the Down’s cell and label the electrodes, electrolytes and the direction of flow of electrons and ions. Write the equations for the reactions involved in the process.
35. Alkali metals are difficult to reduce. Explain why ?
36. Why are alkali metals strong reducing agents ?
37. Reactivity of alkali metals increases with atomic size. Explain.
38. Caesium is the most reactive among alkali metals.Explain why ?
39. An alkali metal has the lowest ionisation energy in a period. Explain.
40. Do alkali metals form only +1 ions ?
41. Why do alkali metals impart characteristic colours to flame ?
42. Explain why alkali metals electrically conducting ?
43. The hydroxides and carbonates of sodium and potassium are easily soluble in water ; while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain.
44. Sodium cannot be obtained by the electrolysis of aqueous sodium chloride solution. Explain.
45. Give two equations of the reactions to indicate that sodium hydroxide is a strong alkali.
46. List the important uses of caustic soda.
47. Discuss the various reactions of Solvay process.
48. Why can potassium carbonate not be prepared by Solvay’s process ?
49. List the important uses of sodium carbonate.
50. How does magnesium occur in nature ? How is magnesium metal obtained by electrolysis method ?
51. Describe briefly the Dow’s process of obtaining magnesium chloride from sea-water.
52. Anhydrous magnesium chloride cannot be obtained by heating hydrated magnesium chloride MgCl2.6H2O. Explain.
53. What is the effect of heat on the following compounds ?
(i) calcium carbonate
(ii) magnesiusulphate heptahydrate
(iii) magnesium chloride hexahydrate
(iv) Gypsum
54. Bones contain calcium ions. What would the ions associated with it ?
55. A piece of burning magnesium ribbon continues to burn in sulphur dioxide. Explain.
56. It is necessary to add gypsum in the final stages of preparation of cement. Explain.
57. What is the function of adding gypsum during final grinding of cement ?
58. Explain why alkali metals form soluble hydroxides, which become stronger alkalies on descending the group.
59. Group 2 elements(Mg and Ca) are harder and denser than group 1 metals (Na and K). Explain.
60. A piece of magnesium ribbon continues to burn in sulphur dioxide. Explain.
61. Write the equation of chemical reaction that takes place when a piece of sodium is put over water. What will happen to the colour of a red litmus paper , if it is dipped in the solution formed in the above reaction ?
62. Caustic alkalies like NaOH are not stored in aluminium vessels. Explain why ?
63. Name the important uses of the following compounds :
(i) Sodium carbonate (iii) Sodium hydroxide
(ii) Epsom salt (iv) Quick lime / Slaked lime
64. Describe an electrolytic method for the extraction of aluminium from its oxide ore.
65. What chemical test would you perform to detect sodium, magnesium calcium and aluminium ions in a given solution ?
66. Describe important properties of sodium carbonate.
67. Describe the preparation and important properties of quick-lime.
68. Give the method of preparation of calcium hydroxide.
69. What is lime water and milk of lime ?
70. Mention the important properties of calcium hydroxide.
71. What is the difference between soda ash and washing soda ?
72. Give the preparation and uses of :
(i) Gypsum (ii) Plaster of paris
73. What are the special properties of plaster of paris which make it useful in hospitals ?
74. What is the difference between quick-lime and slaked lime ? Write their chemical formulae and few reactions.
75. Write a brief account of baking powder.
76. What happens when :
(i) carbon dioxide is passed through slaked lime ?
(ii) sodium carbonate is strongly heated ?
(iii) quick lime is added to water.
77. Give your remarks on :
(a) anhydrous copper sulphate is used for testing the presence of water.
(b) soda ash is different from washing soda.
78. Why anhydrous calcium sulphate is used as a drying agent ?
79. Aqueous solution of sodium carbonate is alkaline. Explain why ?
80. Give the preparation, properties and uses of aluminium chloride.
81. Why is calcium used for removing traces of moisture from alcohol ?
82. What happens when aluminium is heated with :
i) sodium hydroxide ii) Con. Sulphuric acid
83. Name one source in which sodium chloride is found in abundance.
84. Write balanced chemical equation for the reaction when :
i) copper reacts with dilute nitric acid
ii) copper is heated with con sulphuric acid
iii) copper reacts with con HNO3
85. What is the action of H2 on sodium and calcium ?
86. What are alums ? Give examples and uses of alums.
69. Write a note on :
i) glass ii) cement
87. Describe the preparation of potash alum.
88. Name essential substances required for the manufacture of glass.
89. Name the essential substances required for the manufacture of cement.
90. Which property of cement is utilised in the constructional activities ?
91. Describe briefly the manufacture of cement ?
92. Why anhydrous calcium sulphate is used as a drying agent ? Why plaster of paris cannot be employed for this purpose ?
93. Name two elements which soft enough to cut with an ordinary knife.
94. Write three general characteristics of elements of s-block of the periodic table which distinguish them from the elements of other blocks.
95. The alkali metals follow the noble gases in their atomic structure. What properties of these metals can be predicted from this information.
96. Name the alkali metals Which forms super oxide when heated with excess of air.
97. Name the metal which floats on water without any apparent reaction with water.
98. When an alkali metal dissolves in liquid ammonia the solution acquires different colours. Explain the reasons for this type of colur change.
99. What is meant by ‘diagonal relationship’ in the periodic table ? What is it due to ?
100. Describe in detail the manufacture process of sodium carbonate by the Solvay’s process. State the principles involved in the reaction.
101. Describe the industrial uses of caustic soda. Describe one method to manufacture sodium hydroxide. What happens when sodium hydroxide reacts with :
(i) aluminium metal (ii) CO2 (iii) SiO2
102. Why is s-block elements never occur free in nature ? What are their usual modes of occurrence and how they are generally prepared ?
103. Starting from sodium chloride how would you proceed to prepare :
(i) sodium metal (ii) sodium hydroxide
(iii) sodium peroxide (iv) sodium carbonate
104. Explain what happens when :
(i) sodium hydrogen carbonate is heated.
(ii) Sodium amalgam reacts with water
(iii) Fused sodium metal reacts with ammonia
105. Mention the general trends in Grop 1 and Group 2 with increasing atomic number with respect to :
(i) density
(ii) melting point
(iii) atomic size
(iv) ionisation enthalpy
106. How do the following properties change on moving from group 1 to group 2 in the periodic table ?
107. Like Lithium in group 1, beryllium shows anomalous behaviour in group 2. Write three such properties of beryllium which made it anomalous in the group.
108. Comment on each of the following observations :
(i) mobilities of the alkali metal ions in aqueous solution are Li+ < Na+ < K+ < Rb+ < Cs+.
(ii) Lithium is the only alkali metal to form a nitride directly.
(iii) E for M2+(aq) + 2 e M(s) ( where M = Ca, Sr or Ba) is nearly constant.
(iv) LiF is least soluble among fluorides of alkali metals.
109. Compare and constrast the chemistry of Group 1 metals with that of Group 2 with respect to :
(i) nature of oxides.
(ii) Solubility and thermal stability of carbonates
(iii) Polarising power of cations
(iv) Reactivity and reducing power
110. Draw the structures of :
(i) BeCl2 (vapour) (ii) BeCl2 (s)
111. List the raw materials required for the manufacture of portland cement. What is the role of gypsum in it ?
112. How would you explain ?
(i) BeO is insoluble but BeSO4 is insoluble in water.
(ii) BaO is soluble but BaSO4 is insoluble in water
(iii) LiI is more soluble than KI in water.