INTRODUCTION

The elements in the long form of the periodic table has been divided into four blocks, namely s, p, d & f blocks. The elements of group I & II receive their last electron in s-orbital. So they are called as s – block elements.The metals Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr) which have one electron in their outermost shell belongs to group I and are called alkali metals as they react with water to form hydroxides which are strong bases or alkalies. The elements of group II are Berryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), barium (Ba) and radium (Ra) which have two electrons in their outermost shell. All these elements are also metallic in nature and are commonly known as alkaline earth metals with the exception of beryllium. Because of their low density, alkali metals and alkaline earth metals are called lighter metals. Both alkali and alkaline earth metals are highly reactive and hence do not occur in free state but found in combined state. Whereas alkali metals mostly occur as halides, oxides, silicates, borates and nitrates, alkaline earth metals mainly occur as silicates, carbonates, sulphates and phosphates. Some alkali & alkaline earth metals occur abundantly in nature. Calcium is the fifth, magnesium is the sixth, sodium is seventh and potassium is eight barium is the fourteenth and strontium is the fifteenth most abundant element by weight in the earth’s crust. Sodium and magnesium are also present in sea water brine wells and few salt lakes. Anamolous behavour of first element

The first element of a group differs considerably from the rest of the elements of the same group. This anomolous behaviour is due to (i) Smaller size of their atoms (ii) Their higher ionization energies(iii) Their higher electronegativites (iv) Absence of vacant d – orbitals in their valence shell(v) High polarizing power of its cation. Thus Li differes from the rest of alkali metals (Na, K, Rb & Cs) and Be differs from rest of the alkaline earth metals (Mg, Ca, Sr & Ba)

Diagonal relationships

On moving diagonally some members show similar properties with the members of next higher group which is particularly seen in the elements of second and third periods of the periodic table. However the similarities shown are far less pronounced than the similarities with in a group.

The main reasons for the diagonal relationship are

(i) Similarity in electropositive character

The electropositive character decreases along a period, but increases down a group. Hence on moving diagonally the two opposing trend partially cancels out. As a result diagonally related elements have similar electropositive character and hence exhibit similar properties.

1

(ii) Similarity in polarizing power

On moving along a period from left to right, the charge on the ions increases while ionic size decreases, hence polarizing power increases. On moving down the group, the ionic size increases and hence polarizing power decreases. On moving diagonally, these two trends partially cancel out. As a result thus diagonally related elements have same polarizing power and thus exhibit similar properties.

(iii) Similarity in atomic or ionic radii

The atomic and ionic radii decrease across a period and increase in a group. Evidently on moving diagonally, the two trends partially cancel out. As a result, diagonally related elements have similar atomic & ionic radii and hence have similar properties.

Alkali Metals

The group I comprising Li, Na, K, Rb, Cs & Fr are commonly called alkali metals. Francium is radioactive and has a very short life (half life of 21 minutes), therefore very little is known about it.

Electronic Configuration

The general electronic configuration of alkali metals may be represented by [noble gas] where n = 2 to 7

Table – I: [Some Physical properties of group 1 elements (alkali metals)]

Property Elements

Li Na K Rb Cs Fr Radioactive)

Atomic number Electronic configurationAtomic mass Metallic radius (pm) Ionic radius

Ionization enthalpy I II

Electronegativity (Pauling Scale)

Melting point/KBoiling point/KE(V) at 298K for

Occurrence in Atmosphere

3 11 19 37 55 87

6.94 22.99 39.10 85.47 132.91 223152 186 227 248 265 37576 102 138 152 167 180 520 496 419 403 376 - 7298 3562 3051 2633 2230 -

0.98 0.93 0.82 0.82 0.79 0.53 0.97 0.86 1.53 1.90 454 371 336 312 302 1615 1156 1032 961 944

-3.04 -2.714 -2.925 -2.930 -2.927 18* 227** 1.84** 78.12* 2.6*

* ppm (parts per million) ** percentage by weight.

Illustration 1. Why are Group 1 elements called alkali metals?

Solution: The Group 1 elements are called alkali metals because they form water soluble hydroxides.

Exercise 1.Write three general characteristics of the elements of s-block of the periodic table which distinguish them from the elements of the other blocks.

Exercise 2.Which alkali metal is most abundant in earth’s crust?

PHYSICAL PROPERTIES

(i) All the alkali elements are silvery white solid. These are soft in nature and can be cut with the help of knife except the lithium. When freshly cut, they have a bright lusture which quickly fades due to surface oxidation. These are highly malleable and ductile. The silvery luster of alkali metals is due to the presence of highly mobile electrons of the metallic lattice. There being only a single electron per atom, the metallic bonding is not so strong. As the result, the metals are soft in nature. However, the softness increases with increase in atomic number due to continuous decrease in metallic bond strength on account of an increase in atomic size.

(ii) Atomic and Ionic radiiThe atoms of alkali metals have the largest size in their respective periods. The atomic radii increase on moving down the group among the alkali metals.

Reason

On moving down the group a new shell is progressively added. Although, the nuclear charge also increases down the group but the effect of addition of new shells is more predominant due to increasing screening effect of inner filled shell on the valence s-electrons. Hence the atomic size increases in a group. Alkali metals change into positively charged ions by losing their valence electron. The size of cation is smaller than parent atom of alkali metals. However, within the group the ionic radii increase with increases in atomic number. The alkali metal ions get extensively hydrated in aqueous solutions. Smaller the ion more is the extent or degree of hydration. Thus, the ionic radii in aqueous solution follow the order

The charge density on Li+ is higher in comparison to other alkali metals due to which it is extensively hydrated.

(iii) Ionization Energy (Ionization enthalpy)

The first ionization energy of the alkali metals are the lowest as compared to the elements in the other group. The ionization energy of alkali metals decreases down the group.

Reason

The size of alkali metals is largest in their respective period. So the outermost electron experiences less force of attraction from the nucleus and hence can be easily removed. The value of ionization energy decreases down the group because the size of metal increases due to the addition of new shell along with increase in the magnitude of screening effect.

(iv) Oxidation State

The alkali metals show +1 oxidation state. The alkali metals can easily loose their valence electron and change into uni-positive ions

ReasonDue to low ionization energy, the alkali metals can easily lose their valence electron and gain stable noble gas configuration. But the alkali metals cannot form ions as the magnitude of second ionization energy is very high.

(v) Reducing Properties

The alkali metals have low values of reduction potential (as shown in table-I) and therefore have a strong tendency to lose electrons and act as good reducing agents. The reducing character increases from sodium to caesium. However lithium is the strongest reducing agent.

ReasonThe alkali metals have low value of ionization energy which decreases down the group and so can easily lose their valence electron and thus act as good reducing agents.

(vi) Melting and Boiling Points

The melting and boiling points of alkali metals are very low because the intermetallic bonds in them are quite weak. And this decreases with increase in atomic number with increases in atomic size. (vii) DensityThe densities of alkali metals are quite low as compared to other metals. Li, Na and K are even lighter than water. The density increases from Li to Cs.

Reason

Due to their large size, the atoms of alkali metals are less closely packed. Consequently have low density. On going down the group, both the atomic size and atomic mass increase but the increase in atomic mass compensates the bigger atomic size. As a result, the density of alkali metals increases from Li to Cs. Potassium is however lighter than sodium. It is probably due to an unusal increase in atomic size of potassium.

(viii) Nature of bond formed

All the alkali metals form ionic (electrovalent) compounds. The ionic character increases from Li to Cs because the alkali metals have low value of ionization energies which decreases down the group and hence tendency to give electron increases to form electropositive ion.

(ix) Conductivity

The alkali metals are good conductors of heat and electricity. This is due to the presence of loosely held valence electrons which are free to move throughout the metal structure.

(x) Photoelectric Effect

Alkali metals (except Li) exhibit photoelectric effect (A phenomenon of emission of electrons from the surface of metal when light falls on them). The ability to exhibit photoelectric effect is due to low value of ionization energy of alkali metals. Li does not emit photoelectrons due to high value of ionization energy.

(xi) Flame colouration

The alkali metals and their salts impart a characteristic colour to flame

Li Na K RbColour Crimson Red Golden

YellowPale Violet Violet Sky Blue

/nm 670.8 589.2 766.5 780.0 455.5

On heating an alkali metal or its salt (especially chlorides due to its more volatile nature in a flame), the electrons are excited easily to higher energy levels because of absorption of energy. When these electrons return to their ground states, they emit extra energy in form of radiations which fall in the visible region thereby imparting a characteristic colour to the flame.

Illustration 2. What is the most reactive alkali metal and why?

Solution: The most reactive alkali metal is cesium due to its lowest first ionization enthalpy and lowest electronegativity.

Illustration 3. The alkali metals have low densities. Explain.

Solution: The alkali metals have low densities due to their large atomic sizes. In fact, Li, Na and K are even lighter than water.

Illustration 4. The metallic lustre exhibited by sodium is explained by(A) Diffusion of sodium ions(B) Oscillation of loose electrons(C) Excitation of free protons(D) Existence of body–centered cubic lattice

Solution: (B)

Exercise 3.Why is the density of potassium less than sodium?

Exercise 4.Why is lithium the strongest reducing agent in the periodic table?

Exercise 5.Name the metal which floats on water without any apparent reaction with it.

Exercise 6.Which is softer – Na or K and why?

Exercise 7.What makes sodium highly reactive?

Exercise 8.Alkali metals impart colour to Bunsen flame due to(A) The presence of one electron in their outermost orbital(B) Low ionization energies(C) Their softness(D) Their reducing nature

CHEMICAL PROPERTIES

The alkali metals are highly reactive metals and the reactivity increases down the group. The reactivity is due to-(a) low value of first ionization energy (b) large size (c) low heat of atomization

(i) Reaction with OxygenThe alkali metals tarnish in air due to the formation of an oxide or hydroxide on the surface. Alkali metals when burnt in air form different kinds of oxides. For example the alkali metals on reaction with limited quantity of oxygen form normal oxides of formula, M2O

(Where M = Li, Na, K, Rb, Cs) When heated with excess of air, lithium forms normal oxide, ; sodium forms peroxide,, whereas potassium rubidium and caesium form superoxides having general formula

Thus the reactivity of alkali metals with oxygen increases down the group. Further, the increasing stability of peroxide or superoxide, as the size of the metal ion increases is due to the stabilization of larger anions by larger cation through higher lattice energies. Due to small size, has a strong positive field around it which attracts the negative charge so strongly that it does not permit the oxide anion to combine with another oxygen to form peroxide ion, . On the other hand, ion because of its large size than Li+ ion has comparatively weaker positive field around it which cannot prevent ion to combine with another oxygen to form peroxide ion . The larger , , and ions have still weaker

positive field around them which cannot prevent even peroxide ion, to combine with another

oxygen atom to form superoxide .

(ii) Reaction with Hydrogen

Alkali metals react with dry hydrogen at about 673K to form colourless crystalline hydrides. All the alkali metal hydrides are ionic solids with high melting points.

(M = Li, Na, K, Rb or Cs)

Some important features of hydrides are

(a) The stability of hydrides decrease from Li to Cs. It is because of the fact that M-H bond becomes weaker due to increase in the size of alkali metals down the group.

(b) These hydrides react with water to form corresponding hydroxides and hydrogen gas.

(c) These hydrides are strong reducing agents and their reducing nature increases down the group. Alkali metals also form complex hydrides such as and which are good reducing agents.

(d) All these hydrides react with proton donors such as water, alcohols, gaseous ammonia and alkynes liberating gas.

The order of reactivity of the alkali metals towards hydrogen decreases as we move down the group from Li to Cs. This is due to the reason that the lattice energies of these hydrides decreases progressively as the size of the metal cation increases and thus the stability of these hydrides decreases from LiH to CsH.

(iii) Reaction with water

The alkali metals are known to have large negative reduction potential values. As a result they can act as better reducing agents as compared to hydrogen. Hence, alkali metals react with water and other compounds containing acidic hydrogen atoms such as hydrogen halides (HX) and acetylene (C2H2) and liberate H2 gas

The reaction becomes more and more violent as we move down the group. Thus, Lithium reacts gently, sodium melts on the surface of water and the molten metal moves around vigorously and may sometimes catch fire. Potassium melts and always catches fire and so are Rb and Cs.

(iv) Reaction with halogens

Alkali metals react vigorously with halogens to form metal halides of general formula MX, which are ionic crystalline solids.

M = Li, Na, K, Rb or Cs and X = F, Cl, Br or I

Reactivity of alkali metals with particular halogens increases from Li to Cs. On the other hand, reactivity of halogens decreases from to .

(v) Solubility in liquid ammonia

All alkali metal dissolve in liquid ammonia giving deep blue solutions which are conducting in nature. These solutions contain ammoniated cations and ammoniated electrons as shown below:

The blue colour of the solution is considered to be due to ammoniated electrons which absorb energy corresponding to red region of the visible light for the their excitation to higher energy levels. The transmitted light is blue which imparts blue colour to the solutions. The electrical conductivity of the solution is due to both ammoniated cations and ammoniated electrons. The blue solution on standing slowly liberates hydrogen resulting in formation of amide :

At concentrations above 3M, the solutions of alkali metals in liquid ammonia are copper-bronze coloured. These solutions contains clusters of metal ions and hence possess metallic lusture. The blue coloured solutions are paramagnetic due to presence of large number of unpaired electrons, but bronze solutions are diamagnetic due to formation of electron clusters in which ammoniated electrons with opposite spin group together These solutions are stronger reducing agents than hydrogen and hence will react with water to liberate hydrogen.

(vi) Reaction with sulphur and phosphorus

Alkali metals react with sulphur and phosphorus on heating to form sulphides and phosphides respectively.

(vii) Reaction with MercuryAlkali metals combine with mercury to form amalgams. The reactions is highly exothermic in nature

Illustration 5. What happens when (give chemical equation only):(i) sodium is exposed to moist air(ii) sodium reacts with water

Solution: (i)

(ii)

Illustration 6. On prolonged exposure to air, sodium finally changes to(A) Na2CO3 (B) Na2O(C) NaOH (D) NaHCO3

Solution: (B)

Exercise 9.Write balanced equations for reaction between(i) Na2O2 and water(ii) KO2 and water(iii) Na2O and CO2

Exercise. 10.Which reaction below is explosive?(A) Lithium with water (B) Magnesium with water(C) Beryllium with water (D) Caesium with water

General characteristic of the compounds of the alkali metals (i) Oxides and HydroxidesAll the alkali metals, their oxides, peroxides and superoxides readily dissolve in water to produce corresponding hydroxides which are strong alkalies eg

Thus peroxides and superoxides also act as oxidizing agents since they react with forming H2O2 and respectively. The hydroxides of all the alkali metals are white crystalline solids. They are strongest of all base and readily dissolve in water with the evolution of much heat.

(a) Basic Strength

The basic strength of these hydroxides increases as we move down the group Li to Cs. The hydroxides of alkali metals behave as strong bases due to their low ionization energies which

decrease down the group. The decrease in ionization energies leads to weakening of the bond between metal and hydroxide ion and M – O bond in M – O – H can easily break giving and

. This results in the increased concentration of hydroxyl ions in the solution i.e increased

basic characters.

(b) Solubility and stability

All these hydroxides are highly soluble in water and thermally stable except lithium hydroxide.

(c) Formation of salts with acids:

Alkali metals hydroxides being strongly basic react with all acids forming salts.

Illustration 7: Name the alkali metal which forms superoxides when heated in excess of air and why?

Solution: Potassium forms superoxides when heated in excess of air. This is due to the stabilization of large size cation by large size anion.

Illustration 8. Why lithium forms only lithium oxide and not peroxide or superoxides?

Solution: Due to the small size of lithium, it has a strong positive field around it. On combination with the oxide anion (O2–), the positive field of lithium ion restricts the spread of negative charge towards another oxygen atom and thus prevents the formation of higher oxides.

Exercise 11.What happens when potassium superoxide is dissolved in water?

Exercise 12.Account for the fact that the basicity of the alkali metal hydroxides increases down the group.

(ii) Halides

The alkali metals combine directly with halogens under appropriate conditions forming halides of general formula MX. These halides can also be prepared by the action of aqueous halogen acids (HX) on metals oxides, hydroxides or carbonate.

(M = Li, Na, K, Rb or Cs)

(X = F, Cl, Br or I)

All these halides are colourless, high melting crystalline solids having high negative enthalpies of formation.

Table – II [Standard enthalpies of formation in (kJ/mol1)]

Element MF MCl MBr MILi -612 -398 -350 -271Na -569 -400 -360 -288K -563 -428 -392 -328Rb -549 -423 -389 -329Cs -531 -424 -395 -337

The value decreases in the order: Fluoride > Chloride > bromides > IodideThus fluorides are the most stable while iodides are the least stable. The trends in melting points, boiling points and solubility of alkali metals halides can be understood in terms of polarization effects, lattice energy and hydration of ions.

(a) Polarization effects

Comparison of ionic and covalent character of alkali metal halides. When a cation approaches an anion, the electron cloud of the anion is attracted towards the cation and hence gets distorted. This effect is called polarization. The power of the cation to polarize the anion is called its polarizing power and the tendency of the anion to get polarized is called its polarizability. The greater the polarization produced more is the concentration of the electrons between the two atoms thereby decreasing the ionic character or increasing the covalent character. The covalent character of any compound in general depends upon the following factors.

(i) Size of the cations

Smaller the cation greater is its polarizing power and hence larger is the covalent character. The covalent character decreases as size of cation increases.

LiCl > NaCl > KCl > RbCl > CsClThus LiCl is more covalent than KCl.

(ii) Size of the anion

Larger the anion, greater is its polarizability. This explains the covalent character of lithium halides is in order

LiI > LiBr > LiCl > LiF

(iii) Charge of the ion

Greater the charge on the cation greater is its polarizing power and hence larger is the covalent character. The covalent character of some halides increase in the order

Similarly greater the charge on the anion, more easily it gets polarized thereby imparting more covalent character to the compound formed eg covalent character increase in the order

Thus the covalent character decreases as the charge of the anion decrease.

(iv) Electronic configuration of the cation

If two cations have the same charge and size, the one with pseudo noble gas configuration i.e having 18 electrons in the outermost shell has greater polarizing power than a cation with noble gas configuration i.e having 8 electrons. For example CuCl is more covalent than NaCl.

(b) Lattice Energies

lattice energy is defined as the amount of energy required to separate one mole of solid ionic compound into its gaseous ions. Evidently greater the lattice energy, higher is the melting point of the alkali metals halide and lower is its solubility in water Table – III

Compound Lattice energy Hydration* energy Solubility Melting pointLiCl -845 -876 63.7 887NaCl -770 -776 35.7 1084KCl -703 -700 34.7 1039

RbCl -674 -680 77.0 988CsCl -644 -646 162 925NaF -893 -919 4.22 1261NaCl -770 -776 35.7 1028NaBr -730 -745 116 1084NaI -685 -685 184 944LiF -1005 -1019 0.27 1115

CsI -582 -670 44.0 1115

(c) Hydration Energy

It is the amount of energy released when one mole of gaseous ions combine with water to form hydrated ions.

Higher the hydration energy of the ions greater is the solubility of the compound in water. Further the extent of hydration depends upon the size of the ions. Smaller the size of the ion, more highly it is hydrated and hence greater is its hydrated ionic radius and less is its ionic mobility (Conductance). From above arguments, the melting point and solubility in water or organic solvent of alkali metal halides can be explained(i) A delicate balance between lattice enthalpy and hydration enthalpy determines the ultimate

solubility of a compound in water. For eg. Low solubility of LiF (0.27 g/100 g ) is due to its

high lattice energy whereas the low solubility of CsI (44g/100g ) is due

to smaller hydration energy of the two ions (-670 KJ/mol)(ii) The solubility of the most of alkali metal halides except those of fluorides decreases on

descending the group since the decrease in hydration energy is more than the corresponding decrease in the lattice energy.

(iii) Due to small size and high electronegativity, lithium halides except LiF are predominatantly covalent and hence are soluble in covalent solvents such as alcohol, acetone, ethyl acetate, LiCl is also soluble in pyridine. In contrast NaCl being ionic is insoluble in organic solvents.

(iv) Due to high hydration energy of ion, Lithium halides are soluble in water except LiF which is sparingly soluble due to its high lattice energy.

(v) For the same alkali metal the melting point decreases in the order fluoride > chloride > bromide > iodide because for the same alkali metal ion, the lattice energies decreases as the size of the halide ion increases.

(vi) for the same halide ion, the melting point of lithium halides are lower than those of the corresponding sodium halides and thereafter they decrease as we move down the group from Na to Cs.

The low melting point of LiCl (887 K) as compared to NaCl is probably because LiCl is covalent in nature and NaCl is ionic. Illustration 9. Why are alkali metal halides soluble in water?

Solution: Alkali metal halides are soluble in water due to their high ionic character and low lattice energy.

Exercise 13.Why is LiF less soluble in water?

Exercise 14.Why is CsI less soluble in water?

Salts of oxoacids

Since the alkali metals are highly electropositive, therefore their hydroxides are very strong bases and hence they form salts with all oxoacids . They are

generally soluble in water and stable towards heat. The carbonates of alkali metals are

remarkably stable upto 1273 K, above which they first melt and then eventually decompose to form oxides. , however is considerably less stable and decomposes readily.

This is presumably due to large size difference between Li+ and which makes the crystal

lattice unstable.Being strongly basic, alkali metals also form solid bicarbonates. No other metals forms solid bicarbonates though also exists as a solid. Lithium, however does not form solid bicarbonate though it does exist in solution. All the bicarbonate on gentle heating undergo decomposition to form carbonates with the evolution of . All the carbonates and bicarbonates are soluble in water and their solubilities increase rapidly on descending the group. This is due to the reason that lattice energies decrease more rapidly than their hydration energies on moving down the group.

Illustration 10. Complete and balance the following:(i)(ii)

Solution: (i)(ii)

Exercise 15. Which hydroxide is most basic?(A) KOH (B) RbOH(C) NaOH (D) CsOH

Exercise 16. Which of the following alkali metal halides has the lowest lattice energy?(A) LiF (B) NaCl(C) KBr (D) CsI

Anomolous Behaviour of Lithium and its Diagonal Relationship with Magnesium

The properties of lithium are quite different from the properties of other alkali metals. On the other hand, it shows greater resemblance with magnesium, which is diagonally opposite element of it. The main reasons for the anomalous behaviour of lithium as compared to other alkali metals are (i) The extremely small size of lithium atom and its ion. (ii) Greater polarizing power of lithium ion , due to its small size which result in the covalent

character in its compounds. (iii) Least electropositive character and highest ionization energy as compared to other alkali

metals. (iv) Non availability of vacant d-orbitals in the valence shell.The reason for resemblance of properties of lithium with magnesium is that these two elements have almost same polarizing power. The following points illustrate the anomalous properties of lithium and its diagonal relationship with magnesium: (a) The melting point and boiling point of lithium are comparatively high. (b) Lithium is much harder than the other alkali metals. Magnesium is also hard metal. (c) Lithium reacts with oxygen least readily to form normal oxide whereas other alkali metals

form peroxides and superoxides. (d) is weak base. Hydroxides of other alkali metals are strong bases.

(e) Due to their appreciable covalent nature, the halides and alkyls of lithum and magnesium are soluble in organic solvents.

(f) Unlike elements of group 1 but like magnesium. Lithium forms nitride with nitrogen.

(g) LiCl is deliquescent and crystallizes as a hydrate, . Other alkali metals do not form hydrates. also forms hydrate, .

(h) Unlike other alkali metals lithium reacts directly with carbon to form an ionic carbide. Magnesium also forms a similar carbide.

(i) The carbonates, hydroxides and nitrates of lithium as well as magnesium decompose on heating.

The corresponding salts of other alkali metals are stable towards heat.

(j) Lithium nitrate, on heating, decomposes to give lithium oxide, whereas other alkali metals nitrates decomposes to give the corresponding nitrite.

(k) are the only alkali metal salts which are insoluble in water. The corresponding magnesium compounds are also insoluble in water.

(l) Hydrogen carbonates of both lithium and magnesium can not be isolated in solid state. Hydrogen carbonates of other alkali metals can be isolated in solid state.

Illustration 11. Name the chief factor responsible for the anomalous behaviour of lithium.

Solution: The chief factors responsible for the anomalous behaviour of lithium are:(i) its very small size,(ii) high electronegativity,(iii) high ionization enthalpy and(iv) absence of vacant d-atomic orbital in the valence shell

Difficulties encountered during extraction of alkali metals

Alkali metals, can not be extracted from their ores by the usual methods of extraction of metals because of the following difficulties: (i) Alkali metals are strong reducing agents and hence can not be extracted by reduction of their

oxides or chlorides. (ii) Alkali metals being highly electropositive can not be displaced from the aqueous solutions of

their salts by other metals. (iii) Alkali metals can not be isolated by electrolysis of the aqueous solution of their salts since

hydrogen is liberated at the cathode instead of the alkali metal because the discharge potentials of alkali metals are much higher than that of the hydrogen. However, by using mercury as cathode, the alkali metals can be deposited at the cathode but the alkali metals so deposited readily combines with mercury to form an amalgam from which its recovery is very difficult.Therefore in view of above difficulties, only successful method is the electrolysis of their molten (fused) salts usually chlorides.

Extraction of Lithium

Minerals of Lithium

(i) Spodumene, LiAl(SiO3)2 – 6% Lithium(ii) Triphylite (Li, Na)2PO4 (Fe, Mn)3(PO4)2 – 4% Lithium (iii) Petalite LiAl(Si2O5)4 - 2.7 – 3.7% Lithium (iv) Lepidolite (Li, Na, K)2(SiO3)3(FOH)2 - 1.5% Lithium (v) Amblygonite LiAl(PO4)FIt involves the following steps:

1. Preparation of Lithium chloride

The minerals are first of all converted into lithium chloride by any one of the following methods:

(i) Acid treatment method

The finely powdered silicate ore is first heated to about 1373 K to make it more friable and then with at 523 K. The thus formed is cooled, leached with water and then

filtered to remove silica . The filtrate thus obtained is treated with a calculated amount of to precipitate aluminium and iron as carbonates which are filtered off. Excess of is then added to the filtrate to precipitate . This is filtered and dissolved in HCl to

obtain LiCl which is purified by extraction with alcohol.

(ii) Fusion method

The powdered silicate mineral is fused with and the fused mass is extracted with HCl and filtered. The filterate contains chlorides Li, Al, Ca, Na and K whereas silicon is removed as insoluble residue. The filterate is evaporated to dryness and the residue is extracted with pyridine in which only LiCl dissolves. Pyridine is distilled off while LiCl is left behind. The method discussed above may be summed up in the following flow-sheet.

2. Electrolysis of Lithium chlorideA mixture of dry lithium chloride (55%) and potassium chloride (45%) is fused and electrolysed in an electrolytic cell shown in the figure.

Potassium chloride is added to increase the conductivity of lithium chloride and to lower the fusion temperature. The cell is operated at a temperature of about 723 K and voltage of 8-9 volts is applied. As a result of electrolysis, the following reactions take place: At cathode: At anode :

Chlorine gas, a valuable by product liberated at the anode leaves the cell through the exit while molten lithium rises to the surface of the fused electrolytes and collects in the cast iron enclosure surrounding the cathode. The metal thus obtained is 99% pure and is preserved by keeping it wrapped in paraffin wax. It may be noted here that lithium being the lightest metal known (density = 0.534 g ) can not be stored in kerosene oil since it floats on the surface.

Properties of Lithium

(a) Physical Properties

(i) Lithium is a silvery white metal.(i) It is the hardest alkali metal but still is soft enough to be cut with a knife.(ii) Atomic and ionic radii of lithium are the lowest amongst alkali metals.

(b) Chemical Properties

Lithium is highly reactive element. However, among alkali metals lithium is the least reactive.

Lithium is the only alkali metal which combines directly with nitrogen to form lithium nitride. Lithium nitride is ionic and is ruby red. On heating it decomposes to its constituent element. It also reacts with water evolving ammonia.

Reaction with ammonia Like other alkali metals, lithium dissolves is liquid ammonia to form a deep blue solution due to formation of ammoniated electrons.

However, when NH3 gas is passed over molten lithium. Lithium amide is formed

Uses of Lithium (a) Lithium lead alloy (0.05% Li) which is used for making toughened bearings and sheets for

cables. (b) Lithium – Aluminium alloy has great tensile strength and elasticity like that of mild steel. It is

used for air craft construction. (c) Lithium – magnesium alloy (with 14% Li) is extremely tough and corrosion resistant which is

used for armour plate and aerospace components. (d) Lithium is used for refining of metal like copper and nickel as it combines readily with oxygen

and nitrogen and thus removes the last traces of oxygen and nitrogen. (e) Lithium chloride is used in air conditioning plants to regulate the humidity. (f) Lithium aluminium hydride (LiAlH4) is used as a reducing agent in synthetic organic chemistry (g) Lithium carbonate is used in making special variety of glass which is very strong and is

weather proof.

Sodium

Sodium is the 7th most abundant element by weight found in earth’s crust.

Minerals of sodium(i) Rock salt – NaCl (ii) Chile salt petre – NaNO3 (iii) Glauber’s salt – Na2SO4.10H2O (iv) Borax or sodium borate – Na2B4O7.10H2O(v) Albite or soda feldspar – Na2O.Al2O3.6SiO2 or NaAlSi3O8

Extraction of sodium

Sodium is extracted by Down’s process through electrolysis of fused sodium chloride. In this method, sodium is obtained by the electrolysis of mixture of sodium chloride (40%) and calcium chloride (60%) in fused state. The function of calcium chloride is to lower the operating temperature from 1080 K (m.pt. of NaCl) to about 850 K. The main reason for lowering the temperature are:

(i) The melting point of sodium chloride is very high it is very difficult to maintain it in the molten state during electrolysis.

(ii) Sodium is considerably volatile at the temperature needed for the electrolysis and therefore, a part of the metal produced is vapourised.

(iii) Molten sodium gets dispersed in molten chloride to form a metallic fog (colloidal solution) at high temperature.

(iv) Both sodium and chlorine, the two products of the electrolysis, have a corrosive action on the material of the vessel employed for the electrolysis at such a high temperature.

The cell consists of a steel tank lined with fire bricks. A circular graphite anode is placed in the centre of the cell which is surrounded by a cylindrical iron cathode. The anode and cathode are separated by a steel gauge cylinder through which molten sodium chloride can pass but molten sodium cannot. The purpose of using steel gauge is to keep sodium separate from chlorine which would otherwise react with chlorine. The anode is covered by a dome – shaped steel hood which provides the out – let for the escape of chlorine gas. The molten metal liberated at the cathode moves up and flow into the receiver containing kerosene oil. The following reactions take place:

NaCl Na+ + Cl (Ionization) At cathode: Na+ + e Na

At node: Cl Cl + e Cl + Cl Cl2

Properties of Sodium

Physical Properties

1. Sodium is soft, silvery white metal. 2. It is lighter than water, its density is 0.97 g/cm3. 3. It imparts golden yellow flame when introduced into Bunsen flame.

Chemical Properties Sodium is more reactive than lithium.

(i) Action of air and moisture

It is tarnished rapidly on exposure to moist air. First a film of sodium monoxide, Na2O is formed which changes readily into sodium hydroxide by action of moisture and finally into sodium carbonate by the action of CO2 present in air.

or

It is therefore kept under kerosene.

Action of ammonia

Sodium dissolves in liquid ammonia to form blue solution which is a good conductor of electricity.

However, when ammonia is passed through molten sodium, it yield sodamide evolving H2 gas.

Uses of sodium

(i) Sodium is used as a reducing agent in the extraction of boron and silicon. (ii) Liquid sodium or its alloy with potassium is used as coolent in nuclear reactors. (iii) Sodium as well as sodium amalgam are used as reducing agents in organic synthesis. (iv) It is used in Lassaigne’s test for the detection of N, S and halogen in organic compounds. (v) Sodium is used in sodium vapour lamps. (vi) It is largely used for production of artificial rubber, dyes, drugs etc. (vii) Sodium-lead alloy is used for preparation of tetraethyl lead. Pb(C2H5)4 which is used as an

anti-knocking agent in petrol.

Illustration 12. What happens when(i) sodium metal is dropped in water?(ii) sodium metal is heated in free supply of air?(iii) sodium peroxide dissolves in water?

Solution: (i)

(ii)

(iii)

Exercise 17. Explain what happens when(i) sodium hydrogen carbonate is heated.(ii) sodium amalgam reacts with water.(iii) fused sodium metal reacts with ammonia.

Compounds of Alkali Metals

Oxides of Metal

Oxides of Sodium: The possible oxides of Na are Na2O and Na2O2.

Sodium Monoxide

Preparation

It is obtained by burning sodium at 180o C in a limited supply of air or oxygen and distilling off the excess of sodium in vacuum, or by heating sodium peroxide(Na2O2), nitrate(NaNO3) with sodium

2Na + 1/2O2 Na2O Na2O2 + 2Na 2Na2O2NaNO3 + 10 Na 6Na2O + N2

Properties

It dissolves in water violently yielding caustic soda.Na2O + H2O 2NaOH

Sodium Peroxide (Na2O2)

Preparation

It is formed by heating the sodium metal in excess air or oxygen. The air should be moisture free and temperature required is 300o C.

2Na + O2 Na2O2

Properties

(i) It is pale yellow solid becoming white in air due to the formation of a film of sodium hydroxide and carbonate.

(ii) It dissolves in ice - cold water with hydrolysis, yielding hydrogen peroxide, which decomposes into water and oxygen on warming; whereas it gives oxygen and caustic soda with water at room temperature, Na2O2 + 2H2O 2NaOH + H2O2

Na2O2 + 2H2O 2NaOH + O2

(iii) It dissolves in ice-cold dilute mineral acid yielding H2O2.Na2O2 + H2SO4 Na2SO4 + H2O2

Oxides of Potassium

Potassium forms number of oxides namely,K2O,K2O2 and KO2 . Other two oxides can also exist which are K2O3 and KO3

Potassium monoxide (K2O)

It is solid, yellow when hot and white when cold, obtained by heating potassium nitrate with potassium.

2KNO3 + 10K 6K2O + N2

It dissolves in water to give KOH like Na2O

Potassium Superoxide (K2O2)

Controlled oxidation of potassium in excess air or oxygen at 300o C gives mainly K2O2. It gives H2O2 when dissolves in water.

K2O2 + 2H2O 2KOH + H2O2

Potassium dioxide (KO2)

It is made by burning potassium in a good supply of air or oxygen. It is powerful oxidising agent. It reacts with water giving both oxygen and hydrogen peroxide.

2KO2 + 2H2O 2KOH + H2O2 + O2

Illustration 13. Give some important uses of sodium peroxide.

Solution: Important uses of sodium peroxide are:(i) It is used as a bleaching agent.(ii) It is used in the manufacture of sodium perborate, benzoyl peroxide.(iii) It is used for the purification of air in confined spaces such as submarines.

Exercise 18.Potassium metal is commercially prepared by the reduction of molten KCl with metallic sodium at 850oC (1,123 K). This method is based upon the following principle:(A) Sodium is more reactive than potassium at this temperature(B) Potassium, being more volatile, distils off thus shifting the reaction forward(C) Sodium prefers to bind to chloride ions in preference to potassium ions(D) Potassium and sodium form an alloy at this temperature

Hydroxides of Metals (Na, K)

Sodium Hydroxide (NaOH)

When calcium hydroxide is added to sodium carbonate solution, calcium carbonate is precipitated, leaving sodium hydroxide in solution

Na2CO3 + Ca(OH)2 CaCO3 + 2NaOH

Properties

(i) NaOH is stable towards heat but is reduced to metal when heated with carbon 2NaOH + 2C 2Na +2CO + H2

(ii) FeCl3 + 3NaOH Fe(OH)3 + 3NaClNH4Cl + NaOH NaCl + NH3 (pungent smell) + H2O HgCl2 + 2NaOH HgO (yellow powder) + 2NaCl + H2O

Zn(OH)2 + 2NaOH Na2ZnO2 + 2H2OAl2O3 + 2NaOH 2NaAlO2 + H2OSiO2 + 2NaOH Na2SiO3 + H2O3P + 3 NaOH +3H2O PH3 + 3NaH2PO2

2Al + 2 NaOH + 2H2O 3H2 + 2NaAlO2

Uses

(i) It is used in the manufacture of paper, soap and artificial silk.(ii) It is used in petroleum refining.(iii) It is used for mercerizing cotton. (iv) It is used for the preparation of sodium metal and many salts of sodium.

Potassium Hydroxide or Caustic Potash(KOH)

Pure KOH is obtained by adding potassium sulphate to a hot saturated solution of Ba(OH)2.BaSO4 is filtered off and filtrate is evaporated in a silver dish.

K2SO4 + Ba(OH)2 BaSO4 + 2KOH

Properties

(i) In the absorption of carbon dioxide, caustic potash is preferred to caustic soda, since the KHCO3 formed after sufficient absorption is soluble, while NaHCO3 is insoluble and may,therefore choke the tubes.

(ii) As an alkaline reagent KOH is not used while NaOH is used because NaOH is cheaper than KOH.

Carbonates of Metals (Na, K)

Sodium Carbonate (Washing soda) (Na2CO3)

Solvay Process

Sodium carbonate is generally prepared by a process called the ammonia – soda process or solvay 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.

Sodium hydrogen carbonate so formed precipitates out because of the common ion effect caused due to the presence of excess of NaCl. The precipitated NaHCO3 is filtered off and then ignited to get Na2CO3.

Plant Process

The various parts for the manufacture of Na2CO3 by solvay process have been illustrated in figure below.

Plant used for the manufacture of washing soda(i) Ammonia absorber

A 30% solution of brine is saturated with ammonia (from recovery tower) is introduced into absorber tower. Various impurities like calcium and magnesium salts present in the commercial NaCl precipitate out as corresponding insoluble carbonates. (which comes along with ammonia from the ammonia recovery plant)

The ammoniated brine is filtered to remove the precipitated calcium and magnesium carbonate.

(ii) carbonation Tower

It is high tower fitted with perforated plates. Ammoniated brine solution is made to trickle down from the top of the tower while CO2 gas from the lime kiln is admitted from the base of the tower. CO2 rises through the small perforations and its interaction with ammoniated brine results in the formation of insoluble sodium hydrogen carbonate.

(iii) Filtration

The solution containing crystals of NaHCO3 is drawn off from the base of the carbonation tower and filtered to get NaHCO3.

(iv) The NaHCO3 obtained from the above step is heated strongly in kiln to covert it into sodium carbonate (Na2CO3)

The carbon dioxide produced here is sent to carbonation tower.

(v) Ammonia recovery tower

The filtrate, after removal of NaHCO3 contains ammonium salts such as NH4HCO3 and NH4Cl.

The filtrate is mixed with Ca(OH)2 and is heated with steam in ammonia recovery tower.

The mixture of ammonia and CO2 gases is obtained which is used for saturation of brine while calcium chloride is obtained as a by – product.

(vi) Lime kiln

Here limestone is heated at about 1300 K to obtained CO2 and calcium oxide

The CO2 gas goes to the carbonation tower while lime is slaked with water in tank known as slakes to form ca(OH)2. The overall reaction taking place in solvay process is

Flow sheet diagram of Solvay process

Properties

(i) The aqueous solution absorbs CO2 yielding sparingly soluble sodium bicarbonate.Na2CO3 + H2O + CO2 2NaHCO3

(ii) It dissolves in acids with an effervescence of carbondioxide and is causticised by lime to give caustic soda.Na2CO3 + 2HCl 2NaCl + H2O + CO2

Na2CO3 + Ca(OH)2 2NaOH + CaCO3

(iii) Fusion with silica, sodium carbonate yields sodium silicate.Na2CO3 + SiO2 Na2SiO3 + CO2

(iv) Hydrolysis – being a salt of a strong base (NaOH) and weak acid (H2CO3), when dissolved in water sodium carbonate. Undergoes hydrolysis to form an alkaline solution Na2CO3 + 2H2O H2CO3 + 2NaOH

Uses

(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, paints and textile industries. (v) It is used for washing purposes in laundry. (vi) It is used in the manufacture of glass, borax, soap and caustic soda.

Potassium carbonate (K2CO3)

It is also known as pearl ash. It is made by passing CO2 into a conc. solution of the chloride, containing hydrated mangesium carbonate in suspension at 20C when an insoluble potassium hydrogen magnesium carbonate is precipitated. 2KCl + 3(MgCO3. 3H2O) + CO2 2(MgCO3. KHCO34H2O) + MgCl2The precipitate is separated by filtration, and then decomposed either by heating with water under pressure at 140C or by the action of magnesium oxide below 20C.

2(MgCO3. KHCO3. 4H2O) 2MgCO3 + K2CO3 + 9H2O + CO2

2(MgCO3. KHCO3. 4H2O) + MgO 3(MgCO3. 3H2O) + K2CO3

Properties

(i) It is white, deliquescent solid(ii) K2CO3 resembles Na2CO3 in properties, but is more alkaline and more soluble than Na2CO3.

Illustration 14. Sodium carbonate, which is one of the most important products of the chemical industry, is prepared by the Solvay process based on the interaction of sodium chloride with ammonia and carbon dioxide. The reaction yields(A) NH4HCO3 (B) NH4Cl(C) NaHCO3 (D) (NH4)2CO3

Solution: (A)

Bicarbonates of Metals (Na, K)

Sodium Bicarbonate

A concentrated solution of sodium carbonate absorbs CO2 to give sparingly soluble sodium bicarbonate.

Na2CO3 + CO2 + H2O 2NaHCO3

Properties

(i) It is sparingly soluble in water(ii) When heated between 250°C and 300°C, it is converted into pure anhydrous sodium

carbonate which can be used for standardising acids.2NaHCO3 Na2CO3 + H2O + CO2

Potassium bicarbonate

It is made by absorbing CO2 in moist potassium carbonate and then drying the product in a porous plate.

K2CO3 + H2O + CO2 2KHCO3

Properties

KHCO3 resembles NaHCO3, but is much more soluble in water. The solution is strongly alkaline owing to hydrolysis. KHCO3 + H2O KOH + H2CO3

Chlorides of Metal (Na, K)

Sodium Chloride (NaCl)

It is also called common salt occurs abundantly in nature as rock salt or halite. The most abundant source is sea-water where sodium chloride occurs to the extent of 2.6 – 2.9 percent. The sea water is exposed to the sun and air in large shallow pits. The gradual evaporation of water leading to the crystallization of the salt. The purification is done by dissolving the salt in minimum volume of water and filtering, if necessary, to remove insoluble impurities. The solution is then saturated with a current of dry hydrogen chloride whereby crystals of pure sodium chloride separate out.

Properties

(i) NaCl is a colourless crystalline salt, almost insoluble in alcohol and highly soluble in water.(ii) It gives rise to HCl when heated with conc. H2SO4 and Cl2, with MnO2 plus H2SO4.

NaCl + H2SO4  NaHSO4 + HClNaHSO4 + NaCl Na2SO4 + HCl2NaCl + MnO2 + 2H2SO4  MnSO4 + Na2SO4+ 2H2O + Cl2 

Potassium Chloride

KCl is prepared from fused carnallite – nearly pure KCl separates from the melt, leaving fused MgCl2 behind.KCl, MgCl26H2O KCl + MgCl26H2O

PropertiesIt is colourless cubic crystal like solid soluble in water. Its solubility increases almost linearly with temperature.

Sulphates of Metals: (Na, K)

Sodium Sulphate, Na2SO4

The anhydrous salt known as salt cake, is prepared on an industrial scale by heating strongly sodium chloride with conc. sulphuric acid.

NaCl + H2SO4 NaHSO4+ HClNaCl + NaHSO4 Na2SO4 + HCl

Glauber’s salt or hydrated sodium sulphate, Na2SO410H2O is prepared from salt cake by crystallisation from water below 32°C This temperature represents the transition temperature for Na2SO4 and Na2SO4.10H2O.It is colourless salt, crystallising in large monoclinic prisms. It is exceedingly soluble in water.

Potassium Sulphate, K2SO4

It is obtained by strongly heating potassium chloride with conc. H2SO4

KCl + H2SO4  KHSO4 + HClKCl + KHSO4 K2SO4+ HClIt is colourless crystalline salt, m.p. 1070C. It is less soluble in water than sodium sulphate and has no hydrate like the later.

Exercise19.Plaster of paris by (i) giving out water and (ii) uniting with water gives A and B respectively. What are A and B?

Alkaline Earth Metals

The group 2 of the periodic table consists of six metallic elements. They are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra). The name alkaline earth metals was given to magnesium, calcium, barium & strontium since their oxides were alkaline in nature and these oxide remained unaffected by heat or fire and existed in earth.

Occurrence Like alkali metals, alkaline earth metals are also highly reactive and hence do not occur in the free state but are likely distributed in nature in the combined state as silicates, carbonates, sulphates and phosphates.

Minerals Be – Beryl (Be3Al2Si6O18) & Phenacite (Be2SiO4) Mg – Magnesite MgCO3, Dolomite CaMg(CO3)2, Epsomite MgSO4.&H2O Ca – Limestone (CaCO3), fluoropatite [3(Ca3(PO4)3.CaF2], Gypsum (CaSO4.2H2O),

Anhydrite (CaSO4)Sr – Celestite (SrSO4), Strontianite (SrCO3) Br – Barytes (BaSO4)

Electronic ConfigurationThe general electronic configuration of alkaline earth metals is ns2. Be – 1s22s2 Mg – 1s22s2sp63s2 Ca – 1s22s22p63s23p64s2 Sr – [Kr]5s2 Ba – [Xe]6s2 Ra – [Rn]7s2

Physical Properties of Group II elements (i) Atomic and ionic radiiThe atomic radii as well as ionic radii of the members of the family are smaller than the corresponding members of alkali metals.

(ii) Ionization energyThe alkaline earth metal owing to their large size of atoms have fairly low values of ionization energies as compared to the p – block elements. However with in the group, the ionization energy decreases as the atomic number increases. It is because of increase in atomic size due to addition of new shells and increase in the magnitude of screening effect of the electrons in inner shells. Because their (IE)1 is larger than that of their alkali metal neighbours, the group IIA metals trend to the some what less reactive than alkali metals. The general reactivity trend is Ba > Sr > Ca > Mg > Be.

Illustration 15. The 2nd ionization energies of the elements of group I are higher than those of the elements of group II. Explain.

Solution: The 2nd electron in case of alkali metal is to be removed form a cation (unipostive ion) which has already acquired a noble gas configuration whereas in case of alkaline earth metals, the second electron is to be removes fro a cation which is yet to acquire the stable noble gas configuration therefore, removal of 2nd

electron in case of alkaline earth metals requires much less energy than that in case of alkali metals.

There is sharp increase in third ionization energy due to stable inert gas configuration of m+2 ions. This explains the upper limit of +2 oxidation state for the elements.

(iii) Oxidation state

The alkaline earth metal have two electrons in their valence shell and by losing these electrons, these atoms acquire the stable noble gas configuration. Thus, unlike alkali metals, the alkaline earth metals exhibit +2 oxidation state in their compounds.

Illustration 16. The alkaline earth metals shows +2 oxidation state i.e. they always form divalent cations (M2+). Explain.

Solution: If ionization energy were the only factor involved, than group II elements should have formed monovalent ions i.e. Mg+, Ca+ etc rather than Mg2+, Ca+2 etc. This can be explained as follows: (i) The divalent cations of alkaline earth metals acquires stable inert gas

configuration. (ii) The divalent cations results in stronger lattices then monovalent cations and

hence a lot of energy called lattice energy released during formation of divalent cations than monovalent cation which compensates the high second ionization energy.

(iii) The existence of divalent ions in the aqueous solution is due to greater hydration of the divalent ions which counter balance the high value of second ionization energy.

The heat of hydration (hydration energy) of alkaline earth metals are approximately four times higher than alkali metals of comparable size. e.g. Hhyd for Na+ (size 102 pm) = 397 KJmol1 Hhyd for Ca+2 (size 100 pm) = 1650 KJmol1 Larger hydration energy is due to the fact that the alkaline earth metals ions, because of their much larger charge to size ratio, exert a much stronger electrostatic attraction on the oxygen of water molecule.

Lattice energies decreases as atomic number increases

MO MCO3 MF2

Mg -3923 -3178 -2906Ca -3517 -2986 -2610Sr -3312 -2718 -2459Ba -3120 -2614 -2367

(v) Nature of metallic bonding in alkaline earth metals

The alkali metal two electrons are involved in the metallic bonding. Moreover, sizes of alkaline earth metal ions are smaller than those of alkali metal ions. Consequently, stronger metallic bonds are formed which result in the close packing of the atoms. Due to the presence of stronger metallic bonds, alkaline earth metals have

(a) Higher melting points (b) Higher boiling points (c) higher densities (d) Harder than the corresponding alkali metals.

(vi) Density

The alkaline earth metals are denser and harder than the corresponding alkali metals. The atoms of alkaline earth metals have smaller size and are hence held by stronger metallic bonds, as compared to alkali metals. Therefore, they are more closely packed in their crystal lattice which accounts for high density and increased hardness of these elements.

(vii) Characteristic flame colouration

Expect Be & Mg (due to high ionization energy), the alkaline earth metals impart characteristic colour when introduced into flame of a burner. This property is due to the ease of excitation of their valence electrons. When elements or their compounds are introduced to flame, the electron absorbs energy from the flame and gets excited to higher energy levels. When these electrons return to their ground state, they emit absorbed energy in form of visible light having characteristic wavelengths. Depending upon the wavelength of light emitted, different colours are impart to the flame. Salts (generally chlorides) impart characteristic colours to the Bunsen flame.

Ion ColourCa2+ Brick-redSr2+ CrimsonBa2+ Apple greenRa2+ Carmine – red

Illustration 17. Mg forms Mg2+, but Na2+ does not exist. Explain.

Solution: Na metal after the loss of one electron attains a noble gas configuration of neon. Therefore, the removal of second electron is energetically unfavourable. Hence, Na2+ does not exist.

Exercise 20.Which group 2 metal burns with dazzling brilliance in air?

Exercise 21.The second ionization enthalpy of calcium is more than that of the first and yet calcium form CaCl2 not CaCl. Why?

Exercise 22.Which of the alkaline–earth metal ion does not impart any colour to the flame?(A) Mg (B) Ca(C) Sr (D) Ba

Electropositive or Metallic Character

The alkaline earth metals are highly electropositive and hence metallic and their electropositive or metallic character increases down the group. However they are less electropositive or metallic than the alkali metals. It is due to smaller size and higher ionization energies as compared to alkali metals, hence have less tendency to loose electron than those of alkali metals (group I)

Like the alkali metals they also form predominantly ionic compounds but tendency of covalency is greater, particularly with Be and Mg because of their smaller atomic and ionic radii. Be forms compounds which are essentially covalent.

Melting and boiling points

The alkaline earth metals have higher melting and boiling points as compared to those of alkali metals which is attributed to their small size and more close packed crystal lattice as compared to alkali metals and presence of two valence electrons.

Heat of Hydration The heats of hydration of M2+ decreases with an increase in their ionic size and their

values are greater than that of alkali metal ions. Alkaline earth metal ions, because of their larger charge to size ratio, exert a much

stronger electrostatic attraction on the oxygen of water molecule surrounding them. Since the alkaline earth metals (except Be) tend to lose their valence electrons readily,

they act as strong reducing agents as indicated by E0red values. The particularly less negative value for Be arises from the large hydration energy associated with the small size of Be2+ and the relatively large value of heat of sublimation.

Solubility Basic nature of oxides increases down the group but solubilities of sulphates and

carbonates decrease as ionic size increases. The solubility of most salts decreases with increased atomic weight, though

usual trend is reversed with fluorides and hydroxides in this group.

Physical Properties of groups 2 elements (alkaline earth metals)Property Elements

Be Mg Ca Sr Ba RaAtomic number 4 12 20 38 56 88

Atomic mass 9.01 24.31 40.08 87.62 137.33 226.03

Metallic radius/pm 112 160 197 215 222 -

Ionic radius/pm 51 72 100 118 135 148

Ionization enthalpy I(kJ mol1) II

8991757

7371450

5901146

5491064

503965

509979

Enthalpy of hydration of M2+ ions (kJ mol1)

2494 1921 1577 1443 1305 -

Electronegativity (Pauling Scale)

1.57 1.31 1.00 0.95 0.89 0.9

Density/g mol at 298 K 1.85 1.74 1.55 2.63 3.62 5.5

Melting Point/K 1562 924 1124 1062 1002 973

Boiling point /K 2745 1363 1767 1655 2078 (1973) (uncertain)

E(V) at 298 K for M2+(aq) + 2e M(s) 1.97 2.37 2.87 2.89 2.90 2.92

Occurrence in Lithosphere 2* 2.76**

4.6** 384* 390* 1010**

* ppm (parts per million) ** Percentage by weight Reactivity and Electrode potential

All the alkaline earth metals are highly reactive elements since they have a strong tendency to lose the two valences s-electrons to form the corresponding dipositive ions having inert gas configuration. The high reactivity arises due to their low ionization energies and high negative values of their standard electrode potentials. Further, the chemical reactivity of alkaline earth metals increase on moving down the group because the I.E. decreases and electrode potentials become more and more negative with increasing atomic number from Be to Ra. Thus, beryllium is the least reactive while Ba (or Ra) is the most reactive element. Further since the ionization energies of alkaline earth metals are higher and their electrode potential is less negative than the corresponding alkali metals. They are less reactive than corresponding alkali metals.

Reducing Character

The alkaline earth metals are weaker reducing agents than the alkali metals. Like alkali metals, their reducing character also increases down the group. This is due to the reason that the alkaline earth metals have greater tendency to lose electrons so, they act as reducing agent but since their I.E. are higher and their electrode potentials are less negative than the corresponding alkali metals, therefore alkaline earth metals are weaker reducing agents than alkali metals. The sulphates are stable to heat whereas the carbonates decompose to give MO and CO2, the temperature of decomposition increasing from Mg to Ba. BeCO3 is kept in the atmosphere of CO2

to prevent its decomposition. BeCO3 MgCO3 CaCO3 SrCO3 BaCO3

<100C 540C 900C 1290C 1360C

Occurrence and uses of alkaline earth metals

Elements Abundance Main Minerals Uses Beryllium 2.8 ´ 103% First detected in 1798 in the

gemstone beryl and emerald (Be2Al2Si6O18)

Used in corrosion resistant alloys.

Magnesium 2.33%, 7th most abundant element in earth’s crust

Pure Mg first prepared in 1800, named after the magnesia district in Thessaly Greece where large deposits of the mineral are found

When alloyed with Al, Mg is widely used as structural materials because of its high strength, low density and ease in machining.

Calcium 4.15%, 5th most abundant element in earth’s crust.

CaCO3.2H2O obtained in pure form in 1808, calcium is derived from latin word calx, meaning “lime”

As an alloying agent for hardness in aluminium compounds. Calcium is the primary constituent of teeth and bones.

Strontium 0.038% Discovered in 1787 and named after the small town of strontion (Scotland)

SrCO3 is used for the manufacture of glass for colour TV picture tubes.

Barium 0.042% Found in minerals witherite (BaCO3) and barite (BaSO4) after

BaSO4 is used in medicine as

which it is named. a contrast medium for stomach and intestine X – rays

Radium Traces Isolated as chloride in 1898 from the mineral pitchblende

Used in cancer radiotheraphy

Group IIA (Alkaline earth metals) and groups IIB (Zn, Cd, Hg) Mg acts as a bridge element between IIA and IIB.

Sr.No. Properties IIA(Be, Mg, Ca, Sr, Ba, Ra) IIB (Zn, Cd, Hg)1 Electronic

configuration[Inert gas] ns2 [Inert gas] (n – 1)d10ns2

2. Block S – block d – block 3. Oxidation

state +2 +2, mercury also forms dimeric

4. Nature of oxide

BeO is amphoteric, other oxides are basic.

ZnO is amphoteric, CdO and MgO are basic

5. Nature of Halides

Electron – deficient BeX2, others (MX2) are ionic: MgCl2 < CaCl2 < SrCl2 < BaCl2

ZnCl2, CdCl2 are ionic but less than IIA, HgCl2 is covalent.

6. Nature of sulphates

Less soluble in water and solubility decreases down the group BeSO4

> MgSO4 > CaSO4 > SrSO4 > BaSO4

More soluble than IIA

7. Nature of hydroxides

Solubility of hydroxides increases as we move down the group.

Solubility of hydroxides decrease as we move down the group.

8. Nature of sulphides

Soluble ZnS, CdS, HgS insoluble and precipitate in salt analysis.

9. Reactivity Increases as we move down the group Be < Mg < Ca < Sr < Ba

Decreases as we move down the group Zn > Cd > Hg

Difference between alkaline earth metals and alkali metals

Both alkaline earth metals and alkali metals are s – block elements as the last electron enters the ns – orbital. They resemble with each other in some respects but still there are certain dissimilarities in their properties on account of different number of electrons in the valency shell, smaller atomic radii, high ionization potential, higher electro negativity etc.

Properties Alkaline earth metals Alkali metals 1. Electronic Two electrons are present in the One electron is present in the

configuration valency shall. The configuration is ns2 (bivalent)

valency shell. The configuration is ns1

(monovalent) more electropositive

2. Valency Bivalent Monovalent 3. Electropositive

nature Less electropositive More electropositive

4. Hydroxides Weak bases, less soluble and decompose on heating.

Strong bases, highly soluble and stable towards heat.

5. Bicarbonates These are not known in free state. Exist only in solution.

These are known in solid state.

6. Carbonates Insoluble in water. Decompose on heating.

Soluble in water. Do not decompose on heating (LiCO3

is an exception)7. Action of nitrogen Directly combine with nitrogen and

form nitridesDo not directly combine with nitrogen except lithium

8. Action of carbon Directly combine with carbon and form carbides

Do not directly combine with carbon

9. Nitrates Decompose on heating evolving a mixture of NO2 and oxygen

Decompose on heating evolving only oxygen

10. Solubility of salts Sulphates, phosphates fluorides, chromates, oxalates etc are insoluble in water

Sulphates, phosphates, fluorides, chromates, oxides etc are soluble in water.

11. Physical properties Comparatively harder. High melting points. Diamagnetic.

Soft, low melting points paramagnetic.

12. Hydration of compounds

The compounds are extensively hydrated. MgCl2.6H2O, CaCl2.6H2O, BaCl2.2H2O are hydrated chlorides.

The compounds are less hydrated. NaCl, KCl, RbCl form non – hydrated chlorides

13. Reducing power Weaker as ionization potential values are high and oxidation potential values are low.

Stronger as ionization potential values are low and oxidation potential values are high.

Chemical Properties of Group – II elements

Reaction with water – (Formation of hydroxides)

The electrode potential of Be (Be2+/Be = 1.97 V) is least negative amongst all the alkaline earth metals. This means that Be is much less electropositive than other alkaline earth metals and hence does not react with water or steam even at red heat. The electrode potential of Mg (Mg+2/Mg = 2.37 V), although more negative than that of Be yet is still less negative than those of alkali metals and hence it does not react with cold water but reacts with boiling water or steam.

or,

Mg, infact, forms a protective layer of oxide on its surface, therefore, despite its favourable electrode potential it does not react readily with water unless the oxide layer is removed by amalgamating it with mercury. In the formation of oxide film, Mg resembles Al. Ca, Sr and Ba have more negative electrode potentials similar to those of the corresponding group I alkali metals and hence react with even with cold water, liberating H2 and forming the corresponding metal hydroxides. Ca + 2H2O Ca(OH)2 + H2 Reactivity of alkaline earth metals increases as we move down the group. However, the reaction of alkaline earth metals is less vigorous as compared to alkali metals.

Reaction with air (Nitrogen and Oxygen)

(a) Formation of oxides and nitrides

Be metal is relatively unreactive in the massive form and hence does not react below 873K. However, powdered Be is more reactive and burns brilliantly on ignition to give a mixture of BeO & Be3N2.

magnesium is more electropositive than Be and hence burns with dazzling brilliance in air to form a mixture of MgO and magnesium nitride.

Ca, Sr and Ba being even more electropositive react with air readily to form a mixture of their respective oxides and nitrides. The reactivity towards oxygen increases as we go down the group. Thus Ca, Ba and Sr are stored in paraffin but Be and Mg are not because they form protective oxide layer on their surface.

(b) Formation of Nitrides

All the alkaline metals burn in dinitrogen to form ionic nitrides of the formula, M 3N2. This is in contrast to alkali metals where only Li forms Li3N. 3M + N2

Be3N2 being covalent is volatile while the nitrides of all other elements are crystalline solids. All these nitrides decompose on heating and react with water liberating NH3.

(c) Formation of Peroxides

Since larger cations stabilize larger anions. Therefore, tendency to form peroxide increases as the size of the metal ion becomes larger. Thus BaO2 is formed by passing air over heated BaO at 773K.

SrO2 is prepared in similar way but under high pressure and temperature. CaO2 is not formed this way but can be prepared as the hydrate by treating Ca(OH)2 with H2O2 and then dehydrating the product.

Crude MgO2 has been made using H2O2 but peroxide of beryllium is not known. All peroxide are white crystalline ionic solids containing the peroxide ion . Treatment of

peroxide with acids liberates H2O2.

Reaction with hydrogen – (Formation of hydrides)

All the alkaline earth metals except Be combine with hydrogen directly on heating to form metal hydrides of formula MH2.

The hydride of beryllium can also be obtained by the reduction of BeCl2 with LiAlH4.

Both BeH2 and MgH2 are covalent compounds having polymeric structures in which H – atoms between beryllium atoms are held together by three centre – two electron (3C 2e) bonds as shown below:

The hydrides of other elements of this groups i.e. CaH2, SrH2 and BaH2 are ionic and contain the H ions. All the hydrides of alkaline earth metals reacts with water liberating H2 gas and thus act as reducing agents.

CaH2 is called Hydrolith and is used for production of H2 by action of water on it.

Reaction with carbon – (Formation of carbides)

When BeO is heated with carbon at 2175 – 2275 K a brick red coloured carbide of the formula Be2C is formed

.

It is a covalent compound and react water forming methane.

The rest of the alkaline earth metals (Mg, Ca, Sr & Ba) form carbides of the general formula, MC 2

either when the metal is heated with carbon in an electric furnace or when their oxides are heated with carbon.

All these carbides react with water producing acetylene gas.

Reaction with Halogens

The alkaline earth metals react with halogens at elevated temperature to form the halides of the types MX2.Action of Acids

The alkaline earth metals readily react with acids liberating hydrogen.

Reaction with Ammonia

Like alkali metal, the alkaline earth metals dissolve in liquid ammonia to give deep blue black

solution from which ammoniates can be recovered.

Illustration 18. How does the basicity of oxides of group 2 increases down the group?

Solution: The basicity increases down the group

Exercise 23.Why is BeCl2 covalent in nature?

General characteristics of compounds of the Alkaline earth metals

(a) Oxides

The oxides MO are obtained either by heating the metals in oxygen or by thermal decomposition of their carbonates.

Expect BeO all other oxides are extremely stable ionic solids due to their high lattice energies. These have high melting point, have very low vapour pressure, are very good conducts of heat, are chemically inert and act as electrical insulators. Therefore, these oxides are used for lining furnaces and hence used as refractory materials. Due to small size of beryllium ion, BeO is covalent but still has high melting point because of its polymeric nature.

(b) Hydroxides

The hydroxides of Ca, Sr & Ba are obtained either by treating the metal with cold water or by reacting the corresponding oxides with water. The reaction of these oxides with H2O is also sometimes called as slaking.

Be(OH)2 and Mg(OH)2 being insoluble are obtained from suitable metal ion solutions by precipitation with OH ions.

Properties

(i) Basic Character All the alkaline earth metal hydroxides are bases except Be(OH)2 which is amphoteric. This basic strength increases as we move down the group. This is because of increase in size which results in decrease of ionization energy which weakens the strength of M – O bonds in MOH and thus increase the basic strength. However, these hydroxides are less basic than the corresponding alkali metal hydroxides because of higher ionization energies, smaller ionic sizes and greater lattice energies.

(ii) Solubility in Water

Alkaline earth metals hydroxides are less soluble in water as compared to alkali metals. The solubility of the alkaline earth metal hydroxides in water increases with increase in atomic number down the group. This is due to the fact that the lattice energy decreases down the group due to increase in size of the alkaline earth metals cation whereas the hydration energy of the cation remains almost unchanged. The resultant of two effects i.e.

becomes more negative as we move from Be(OH)2 to Ba(OH)2 which accounts for increase in solubility.

Halides

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

Properties

(1) All beryllium halides are essentially covalent and are soluble in organic solvents. They are hydroscopic and fume in air due to hydrolysis. On hydrolysis, they produce acidic solution.

(2) The halides of all other alkaline earth metals are ionic. Their ionic character, however increases as the size of the metal ion increase.

(3) Except BeCl2 all other chlorides of group 2 form hydrates but their tendency to form hydrates decreases for eg – MgCl2.6H2O, CaCl2.6H2O.

(4) The hydrated chloride, bromides and iodides of Ca, Sr and Ba can be dehydrated on heating but those of Be and Mg undergo hydrolysis.

(5) BeF2 is very soluble in water due to the high hydration energy of the small Be+2 ion. The other fluorides (MgF2, CaF2, SrF2 and BaF2) are almost insoluble in water. Since on descending the group lattice energy decreases more rapidly than the hydration energy. Therefore whatever little solubility these fluorides have that increase down the group. The chlorides, bromides and iodides of all other elements i.e. Mg, Ca, Sr, Ba are ionic have much lower melting points than the fluorides and are readily soluble in water. The solubility decreases some what with increasing atomic number.

(6) Except of BeCl2 and MgCl2, the other chlorides of alkaline earth metals impart characteristics colour to flame. CaCl2 = Brick red colour SrCl2 = Crimson colour BaCl2 = Grassy green colour

Uses

(i) Calcium fluoride or fluorospar (CaF2) is by far the most important of all the fluorides of the alkaline earth metals since it is the only large scale source of fluorine.

(ii) CaCl2 is widely used for melting ice on roads, particularly in very cold countries because 30% eutectic mixture of CaCl2/ice freezes at 218 K as compared to NaCl /ice at 255K.

(iii) CaCl2 is also used as a desiccant (drying agent) in the laboratory. (iv) Anhydrous MgCl2 is used in the electrolytic extraction of magnesium.

Solubility and Thermal stability of oxo salts

The salts containing one or more atoms of oxygen such as oxides, hydroxides, carbonates, bicarbonates, nitrites, nitrates, sulphates, oxalates and phosphates are called oxo salts.

Sulphates

The sulphates of alkaline earth metals (MSO4) are prepared by the action of sulphuric acid on metals, metals oxides, hydroxides and carbonates.

Properties of sulphates

The sulphates of alkaline earth metals are all white solids.

(a) Solubility

The solubility of the sulphates in water decreases down the groups i.e. Be > Mg > Ca > Sr > Ba. Thus BeSO4 and MgSO4 are highly soluble, CaSO4 is sparingly soluble but the sulphates of Sr, Ba and Ra are virtually insoluble.

Reason

The magnitude of the lattice energy remains almost constant as the sulphate is so big that small increase in the size of the cation from Be to Ba does not make any difference. However the hydration energy decreases from Be+2 to Ba+2 appreciably as the size of the cation increase down the group. Hence, the solubilities of sulphates of alkaline earth metals decrease down the group

mainly due to the decreasing hydration energies from Be+2 to Ba+2. The high solubility of BeSo4

and MgSO4 is due to high hydration energies due to smaller Be+2 and Mg+2 ions.

(b) Stability

The sulphates of alkaline earth metal decomposes on heating giving the oxides and SO3.

The temperature of decomposition of these sulpahtes increases as the basicity of the hydroxide of the corresponding metal increase down the group Carbonates and Bicarbonates

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

hydroxides.

(ii) Sodium or ammonium carbonate is added to the solution of the alkaline earth metal salt such as CaCl2.

Properties All carbonates are stable but beryllium carbonate is prone to hydrolysis. It contains the hydrated

ion rather than Be+2 and hence is precipitated only in an atmosphere of CO2.

Solubility

The carbonates of magnesium and other alkali earth metals are sparingly soluble is water and their solubility decreases down the group from Be to Ba. For e.g MgCO3 is slightly soluble in water but BaCO3 is almost insoluble. The solubility can be explained by the reason same as for sulphates. All the carbonates of alkaline earth metals are however, more soluble in the presence of CO2 due to the formation of corresponding bicarbonates. For e.g.

Stability

The carbonate of all alkaline earth metals decompose on heating to form the corresponding metal oxide and CO2.

The temperature of decomposition, i.e thermal stability of these carbonates, however increase down the group from Be to Ba as the basicity of metal hydroxide increases from Be(OH) 2 to Ba(OH)2. The sulphates are stable to heat whereas the carbonates decompose to give MO and CO2. Thus BeCO3 is unstable and kept in the atmosphere of CO2 to prevent its decomposition.

BeCO3 MgCO3 CaCO3 SrCO3 BaCO3

<100C 540C 900C 1290C 1360C

Bicarbonates

The bicarbonates of alkaline earth metals are prepared by passing CO2 through a suspension of metal carbonates in water.

All the bicarbonates of alkaline earth metals are stable only in solution and have not been isolated in the pure state. Nitrates

Alkaline earth metals nitrates are prepared in solution and can be crystallized as hydrated salts by the action of HNO3 on oxides hydroxides and carbonates.

(M = Be, Mg, Ca, Sr or Ba) Magnesium nitrate crystallizes as Mg(NO3)2.6H2O White Ba(NO3)2 crystallises as unhydrous salt.All nitrates on heating give the corresponding oxides.

(M = Be, Mg, Ca, Sr or Ba)

Illustration19. Arrange the following in order of the property indicated:(i) carbonates of group 2 order of increasing thermal stability(ii) sulphates of group 2 order of decreasing solubility

Solution. (i) BeCO3 < MgCO3 < CaCO3 < SrCO3 < BrCO3 (ii)

Exercise 24.Why is BeCO3 stored in carbon dioxide atmosphere?

Anamolous Behaviour of Beryllium

The properties of berrylium the first member of the alkaline earth metal, differ from the rest of the member. Its is mainly because of (i) Its small size and high polarizing power. (ii) Relatively high electronegativity and ionization energy as compared to other members. (iii) Absence of vacant d – orbitals in its valence shell. Some important points of difference between beryllium and other members (especially magnesium) are given below: (i) Be is harder than other members of its group. (ii) Be is lighter than Mg. (iii) Its melting and boiling points are higher than those of Mg & other members.(iv) Be does not react with water while Mg reacts with boiling water. (v) BeO is amphoteric while MgO is weakly basic. (vi) Be forms covalent compounds whereas other members form ionic compounds.

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

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

Resemblance of Beryllium with Aluminium (Diagonal relationship)

The following points illustrate the anomalous behaviour of Be and its resemblance with Al. (i) Unlike groups – 2 elements but like aluminium, beryllium forms covalent compounds. (ii) the hydroxides of Be, [Be(OH)2] and aluminium [Al(OH)3] are amphoteric in nature, whereas

those of other elements of group – 2 are basic in nature. (iii) the oxides of both Be and Al i.e. BeO and Al2O3 are high melting insoluble solids. (iv) BeCl2 and AlCl3 have bridged chloride polymeric structure.

(v) The salts of beryllium as well as aluminium are extensively hydrolysed. (vi) Carbides of both the metal reacts with water liberating methane gas.

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

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

Illustration 20. Give the structure of BeCl2 in the(i) vapour state (ii) solid state

Solution: (i) Linear molecule(ii) Polymeric structure with bridged chlorine atom

Exercise 25.An aqueous solution of beryllium chloride is acidic in nature. Explain.

Magnesium Metal

Magnesium occurs as magnesite MgCO3, dolomite CaMg(CO3)2, Epsomite (MgSO4.7H2O) and carnalite K2MgCl4.6H2O and langbeinite K2Mg2(SO4)3 deposits. The chloride and sulphate of magnesium occurs in sea water from which it being extracted on an increasing scale.

Extraction

(a) From magnesite or Dolomite

The ore is first calcined to form the oxide MgCO3 MgO + CO2 CaCO3.MgCO3 CaO.MgO + 2CO2

The metal is obtained from the oxide or the mixed oxides as follows: (i) From MgO:

The oxide is mixed with carbon and heated in a current of chlorine gas. MgO + C + Cl2 MgCl2 + CO The chloride thus obtained is subjected to electrolysis.

(ii) The mixed oxides [CaO.MgO] obtained from calcination of Dolomite [CaCO3.MgCO3] are redcued by ferrosilicon under reduced pressure above 1273 K. 2CaO + 2MgO + FeSi 2Mg + Fe + Ca2SiO4

(b) From Carnallite

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

(c) From Sea water

Sea water containing magnesium chloride is concentrated under the sun and is treated with calcium hydroxide Ca(OH)2. Mg(OH)2 is thus precipitated, filtered and heated to give the oxide. The oxide so obtained is treated as in (a) (i) above and then electrolysed.

Electrolysis of Magnesium Chloride

MgCl2 obtained by any of the above methods is fused and mixed with additional mixture of NaCl and CaCl2 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: Mg+2 + 2e Mg At cathode: 2Cl Cl2 + 2e

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

Properties of Magnesium

Physical Properties

(i) Magnesium is a silvery white metal which soon becomes dull in air. (ii) It is a light metal with a density of 1.74 g cm3. (iii) It is fairly malleable and ductile.

Chemical Properties

(i) Action of oxygen or air

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

(ii) Action of CO2 and SO2

Because of its great affinity for oxygen magnesium keeps on burning even in CO2 or SO2.

(iii) Action of nitrogen

On heating magnesium combines with nitrogen to form magnesium nitride.

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

(iv) Action of halogens

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

(v) Action of water

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

(vi) Action of Acids

Dilute acids reacts with magnesium to produce dihydrogen.

However with conc. H2SO4, SO2 is produced

(vii) Reaction with alkyl halide

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

Uses of Magnesium

(i) The chief use of magnesium is in the preparation of alloys with aluminium, zinc, manganese and tin. Duralium (Al = 95%, Cu = 4%. Mn = 0.5%, Mg = 0.5%) Mangnalium (Al = 90% & Mg = 10%) Duralium being light, tough and durable is used for the manufacture of airplanes and automobiles parts. Magnalium being light, tough and hard is used for making balance beams.

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

(iii) Mg is used for ignition of thermite charge in aluminothermy. (iv) A suspension of magnesium hydroxide known as milk of magnesium is used as an antacid

for patients suffering from acidity. (v) In preparation of Grignard reagent. (vi) Being a reducing agent, magnesium is used in the extraction of boron and silicon from their

respective oxides.

Illustration 21. The chemistry of Lithium is very much similar to that of magnesium even though they are placed in different groups - Explain.

Solution: The ratio of their charge to size is nearly same by which they show the diagonal relationship.

Illustration 22. Name some minerals of magnesium.

Solution: Minerals of magnesium are:(i) Magnesite, (ii) Dolomite(iii) Epsom salt

Illustration 23. Write the chemical formulae of the following ores:(i) Dolomite (ii) Epsom salt(iii) Carnallite (iv) Magnesite

Solution: (i) Dolomite – MgCO3.CaCO3

(ii) Epsom salt – MgSO4.7H2O(iii) Carnallite – KCl.MgCl2.6H2O(iv) Magnesite – MgCO3

Exercise 26.What is milk of magnesia? Give its one use.

Exercise 27.

What happens when magnesium reacts with (i) CO2 (ii) SO2 gas?

Compounds of alkaline earth metals

Magnesium sulphate, Epsom salt MgSO4.7H2O

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

Preparation

(i) From dolomite

The dolomite ore is boiled with dil. H2SO4. CaCO3.MgCO3 + 2H2SO4 CaSO4 + MgSO4 + 2H2O + 2CO2 The ppt of calcium sulphate are filtered off and the solution on concentration and cooling gives crystals of MgSO4.7H2O.

(ii) From Magnesite

The magensite ore is powdered and dissolved in dilute H2SO4. The resulting solution is concentrated and cooled when crystals of MgSO4.7H2O separate out. MgCO3 + H2SO4 MgSO4 + H2O + CO2

(iii) From Kieserite

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

(iv) Laboratory Preparation

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

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

Properties

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

(ii) Isomorphism

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

(iii) Action of Heat

When heated it losses 6 molecules of water to give Magnesium sulphate monohydrate which becomes anhydrous when heated to 503 K and finally decomposes to MgO & SO 3 gas on strong heating.

Uses

(i) MgSO4 is used as purgative medicine. (ii) It is used as mordant for cotton in dyeing industry. (iii) It is used in preparation of fire proof textile and wood. (iv) Anhydrous MgSO4 is used as a drying agent in organic chemistry. (v) It is used in preparation of platinised asbestors which is used as a catalyst in the contact

process for the manufacture of H2SO4.

Illustration 24. What are isomorphous salts?

Solution: The salts having the similar crystal structure are called isomorphous salts.Examples: MgSO4.7H2O, ZnSO4.7H2O, FeSO4.7H2O

Oxides of Mg, Ca

MgO(Magnesia)

It is made by heating magnesite (MgCO3).MgCO3 MgO + CO2

It is very slightly soluble in water imparting an alkaline reaction to the solution.MgO + H2O Mg(OH)2

Calcium oxide, Quick lime CaO

Preparation

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

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

Properties

(i) It is a white amorphous solid with m.p. of 2870 K. (ii) When exposed to atmosphere, it absorbs moisture and CO2 forming slaked lime and calcium

carbonate respectively. CaO + H2O (Moisture) Ca(OH)2 CaO + CO2 CaCO3

(iii) On adding water, it produces a hissing sound a large amount of heat is evolved which converts water into steam. This process is called slaking of lime and the fine powder obtained is called slaked lime. CaO + H2O Ca(OH)2 ; H = 63KJ

(iv) Action of acids and acidic oxides

It is a basic oxide and hence combines with acids and acidic oxides forming salts. CaO + 2HCl CaCl2 + H2OCaO + SO2 CaSO3

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

(vi) Reaction with ammonium salt On heating with ammonia salts, it liberates ammonia gas. CaO + 2NH4Cl CaCl2 + 2NH3 + H2O

Uses

(i) It is used as a building materials. (ii) It is used for drying alcohols and non acidic gases. (iii) It is used in the preparation of ammonia and soda lime (CaO + NaOH). (iv) It is used as a basic lining in furnaces.

Illustration 25. Why does a piece of burning magnesium continue to burn in SO2?

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

Exercise 28.Name the substance used for drying alcohol and non-acidic gases.

Hydroxides of Mg & Ca

Magnesium Hydroxide [Mg(OH)2]

It is obtained by adding caustic soda solution to a solution of magnesium sulphate or chloride.MgSO4 + 2NaOH Na2SO4 + Mg(OH)2

Properties

(i) It is converted into its oxide on heating. Mg(OH)2 MgO + H2O

(ii) It dissolves in NH4Cl solution easily.Mg(OH)2 + 2NH4Cl MgCl2 + 2NH4OH

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

Preparation

(i) From Quick lime

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

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

(ii) From calcium chloride

It is obtained by treating calcium chloride with caustic soda.

Properties

Physical

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

Chemical properties

(i) Action of heat

it loses water only at temperature above 700 K.

(ii) Reaction with chlorine

It forms calcium hypochlorite a constituent of bleaching power.

(iii) Reaction with carbon dioxide

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

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

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

(iv) Reaction with acids

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

However, Ca(OH)2 does not dissolve in dil. H2SO4 because the CaSO4 formed is sparingly soluble in water.

Uses

(i) Calcium hydroxide is used for absorbing acidic gases such as CO2, NO2, SO2, SO3 etc. (ii) For preparing ammonia from ammonium salts. (ii) For softening of hard water. (iv) In the laboratory, as lime water for detection of CO2. (v) In white washing due to its disinfectant properties. (vi) In the production of mortar which is used as a building material.

Illustration 26. Which is the weakest base among NaOH, Ca(OH)2, KOH and Be(OH)2

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

Calcium Carbonate (CaCO3)

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

CaCl2 + (NH4)2CO3 CaCO3 + 2NH4Cl

Properties

It dissolves in water containing CO2, forming Ca(HCO3)2 but is precipitated from solution by boiling.CaCO3 + H2O + CO2 Ca(HCO3)2

Illustration 27. Thermal decomposition of a compound 'X' yields, a basic oxide ( Y ) and acidic oxide( Z ) simultaneously. The acidic oxide(Z) can be absorbed by alkaline KOH. What is X,Y,Z.

Solution:

CO2 + 2KOH K2CO3 +H2O

Magnesium Carbonate (MgCO3)

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

MgCl2 + NaHCO3  MgCO3 + NaCl + HCl

Properties

(i) It is very much more soluble in water.(ii) It dissolves in water containing CO2 due to formation of soluble bicarbonate.

MgCO3 + H2O + CO2  Mg(HCO3)2

Bicarbonates of Mg & Ca Calcium bicarbonate [Ca(HCO3)2]

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

Magnesium bicarbonate [Mg(HCO3)2]

Same as in Ca(HCO3)2

Illustration 28. NaHCO3 and NaOH cannot exist together in solution- Why?

Solution: NaHCO3 is an acid salt which must react with NaOH which is strong base. The reaction is as follows:NaHCO3 + NaOH Na2CO3 + H2O

Halides of Mg & Ca

Calcium Chloride (CaCl26H2O)

It separates out as deliquescent crystals when a solution of lime or calcium carbonate in HCl is evaporated.CaCO3 + 2HCl CaCl2 + H2CO3

But it separates out from the reaction mixture as CaCl26H2O. The anhydrous salt is obtained on heating above 200C.

Properties

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

Exercise 29.What is being used as a laboratory dessicant?

Magnesium chloride (MgCl26H2O)

It is prepared in the laboratory by crystallizing a solution of the oxide, hydroxide or carbonate in dilute hydrochloric acid.MgO + 2HCl MgCl2 + H2O

Properties

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

Illustration 29. Complete the following reactions:(i)(ii)

Solution: (i)(ii)

Exercise30.Which magnesium compounds are the constituents of toothpaste?

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

It occurs in nature as gypsum and the anhydrous salt as anhydride. It is prepared by precipitating a solution of calcium chloride or nitrate with dilute sulphuric acid.The effect of heat on gypsum or the dihydrate presents a review of interesting changes. On heating the monoclinic gypsum is first converted into orthorhombic form without loss of water. When the temperature reaches 120C, the hemihydrate or plaster of paris is the product. The latter losses water, becomes anhydrous above 200C and finally above 400C, it decomposes into calcium oxide.

2CaSO4 2CaO + 2SO2 + O2

The following conditions are necessary

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

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

Properties

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

Uses

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

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

Industrial uses of lime and Limestone

Uses of lime

Calcium oxide is called lime or quick lime. It main industrial uses are (i) It is used in steel industry to remove phosphates and silicates as slag. (ii) It is used to make cement by mixing it with silica, alumina or clay. (iii) It is used in making glass. (iv) It is used in lime soda process for the conversion of Na2CO3 to NaOH & vice versa. (v) It is used for softening water, for making slaked lime Ca(OH)2 by treatment with water and

calcium carbide CaC2.

Uses of Slaked lime [Ca(OH)2]

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

(ii) In manufacture of bleaching powder by passing Cl2 gas. (iii) In making glass and in the purification of sugar and coal gas. (iv) It is used in softening of hard water.

Uses of lime stone (CaCO3)

(i) It is used as building material in form of marble. (ii) In manufacture of quick lime. (iii) It is used as a raw material for the manufacture of Na2Co3 in solvay – ammonia process. (iv) Commercial limestone contains iron oxide, alumina, magnesia, silica & sulphur with a CaO

content of 22 – 56% MgO content upto 21%. It is used as such as a fertilizer. Cement

Cement is essentially a finely powdered mixture of calcium silicates and aluminates along with small quantities of gypsum which sets into a hard stone like mass when treated with water. The chief compounds of cement are tricalcium silicate 3CaO.SiO2, dicalcium silicate, 2CaO.SiO2

and tricalcium aluminate 3Ca. Al2O3. Out of these tricalcium silicate is the most important since it has property of setting quickly and acquiring considerable strength within a few days. It usually constitute 50% of the cement.

Composition of Portland cementLime (CaO) - 50 – 60% MgO - 2 – 3% Silica (SiO2) - 20 – 25% Ferric oxide (Fe2O3) - 1 – 2% Alumina (Al2O3) - 5 – 10% Sulphur trioxide (SO3) 1 – 2% For a good quality cement, the ratio of alumina Al2O3 to silica (SiO2) should lie between 2.5 & 4 while that of lime CaO to silica + alumina + ferric oxide should be as close to 2 as possible. Manufacture of cement

Illustration 30. Differentiate between (i) quick-lime (ii) slaked lime and, (iii) lime water

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

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

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

p-Block Elements

Introduction

The right side of the periodic table having group number 13, 14, 15, 16, 17 and 18 are known as p – block elements. These elements have 3, 4, 5, 6, 7 and 8 electrons in their outer most shell, respectively. The last electron of these groups’ elements occupies the position in p – sub shell that is why these elements are called as p – block elements. Their general electronic configuration is ns2np1-6.

Some important properties of p-block(1) Electron affinity

Electron affinity increases from left to right along the period amongst the p – block elements and it decreases from top to bottom. But group 15 is having exceptionally low values of electron affinity and is due to extra stability of exactly half filled orbitals in their valence shell. Similarly,

elements of group 18 (noble gases) have zero affinities due to presence of complete octet which provides them stability. (2) Metallic Character

The metallic character is governed by(i) Size of atoms and (ii) Ionization energy. The elements having bigger size and low ionization energy has a greater metallic character. After combining both the above mentioned factors we observe that the elements with above two properties are located in left corner of p – block and strong non – metallic elements are located at right corner and a diagonal strip of elements separates thus two, having in between properties are called as metalloids.

(3) Oxidation state

The p-block elements show variety of oxidation states both positive and negative. Some of the p-block elements show different oxidation state due to inert – pair effect, where their lower oxidation state is more predominant.

(4) Diagonal relationship

On moving diagonally across the periodic table the element shows certain similarities

Note:Elements of 2nd period differ from their own group elements in some of the properties. This is due to the following reason:(a) Small size(b) Absence of vacant d-orbital(c) High IP

Some important characteristics of p – block in tabular form

Sr. No. Property Along period (left to right) Along group (top to bottom)1. Atomic radii Decreases Increases 2. Ionization potential Increases Decreases3. Electron affinity Increases Decreases4. Electro negativity Increases Decreases5. Metallic character Decreases Increases 6. Oxidizing property Increases Decreases7. Reducing property Decreases Increases

Boron FamilyGroup 13 elements are boron (B), aluminium (Al), Gallium (Ga), Indium (In) and thallium (Tl).Boron is the only non – metal in this group others are metal. Non – metallic character of boron is due to its small size and high ionization energy. The general valence shell electronic configuration of these elements is ns2np1.

General Introduction of Boron family

Boron

The name boron comes form the Arabic and Persian words for borax, its principal ore. It was first isolated in 1808 by Gay – Lussac and Thenard and independently by Sir Hymphry Davy. The pure element is shiny and black. It is very hard and in extremely pure form is nearly as hard as diamond, but too much brittle for practical use. At high temperatures it is a good conductor but at room temperature and below is an insulator.

Aluminium

Aluminium ranks third on the list of the ten most abundant elements in the earth’s crust, while its oxide is fourth among the ten most common compounds in the crust. It is the most abundant metal on the planet. Its name is taken from the Latin alumen for alum. It is soft, light weight and silvery, its existence was proposed by Lavoisier in 1787, it was named by Davy in 1807 and finally isolated by Orsted in 1825. In its purest form the metal is bluish – white and very ductile. It is an excellent conductor of heat and electricity and finds use in some wiring. When pure it is too soft for construction purposes but addition of small amounts of silicon and iron hardened it significantly. Aluminium is the most abundant element in earth crust among this family

Gallium

Gallium is one of the elements originally predicted by Mendeleev in 1871 as aluminium, indicating that it should have the properties similar to aluminium. The actual metal was isolated and named by Paul – Emile Lecog de Boisbaudran in 1875. At room temperature gallium is soft as lead and can be cut with a knife. Its melting point is abnormally low and it will begin to melt in the palm of a warm hand. Gallium is from one of the small numbers of metals that expands on freezing.

Indium

The element indium (Latin indicum, for the colour indigo) was discovered in 1863 by Reich and Ritcher. It is a rare metal, with an abundance similar to that of silver. It is generally found in deposits with zinc and refineries which produce this more common metal often sell indium as well. The pure metal is so soft that you can “wipe” it onto other material in much the same way as lead. It is corrosion resistant.

Thallium

Sir William Crookes discovered thallium in 1861, positively identifying it by a green line in its spectrum (hence the name, which is from the Greek, thallos for “green twig”)Thallium compounds are quite toxic and some have been used as rat poisons. A few compounds are used in glasses for special infra – red lenses.

General Trends in Physical properties

Sr. No. Property Boron Aluminium Gallium Indium Thallium 1. Configuration [He]2s22p2 [Ne]3s2sp2 [Ar]4s24p2 [Kr]5s25p2 [Xe]6s26p2

2. Common oxidation state

+3 +3 +3 +3 +3, +1

3. Atomic radius (pm)

83 143 135 167 170

4. First ionization 801 578 579 558 589

energy (KJ/mol)5. Electro negativity 2.0 1.5 1.6 1.7 1.8

1. Density Generally increases down the group but aluminium has an exceptionally low density.

2. Melting point and Boiling pointB to Ga decrease then Ga to Tl increases Ga has lowest M.P (29.8C) and therefore liquid at room temperature.

3. Atomic radii and Ionic radiiOn moving from B to Tl the size increases due to addition of new energy shells at each step down the group but Ga is smaller than Al.

4. Ionization Energy Generally IE decreases down the group but Ga has higher IE than Al exceptionally due to smaller in size as compared to Al.

5. Metallic character Electropositive character increases down the group hence metallic character also increase down the group but aluminium is having high metallic character than Gallium due to low IE than Ga.

6. Oxidation stateThe elements of B group have three valence electron i.e .two in s – subshell. The most oxidation state should be +3 but due to small size of boron it can not lose its valence electrons to B3+ ion, and combines with other atoms through covalent bonds. Except boron, other elements also exhibit +1 oxidation state and down the group +1 state becomes more stable.

7. Reducing property Down the group the reducing property decreases. Al > Ga > In > Tl

Group trends in chemical properties

1. Hydrides

None of the element from 13 groups reacts directly to hydrogen. However a number of hydrides of these elements have been prepared by indirect methods. Boron hydrides are called boranes. Two types of boranes: (a) BnHn+4, called nidoboranes (b) BnHn+6, called arachnoboranes

The simplest borone is diborane B2H6

Other elements of this group forms only a few stable hydrides of NH 3 types. AlH3 is colour less polymeric solid of formula (AlH3)x and contains Al H Al bridges. A complex hydride of aluminium is a very good reducing agent and used as a regent in lab is Li[AlH 4], Lithium aluminium hydrides. It is a white crystalline solid.

Gallium also form Li[GaH4]

2. Oxides and hydroxides The 13 group elements forms oxides and hydroxides of composition M2O3 and M(OH)3

respectively. As we move down the group the acidic character in oxides and hydroxides

decreases and basic character increases due to decrease in strength of M O bond due to which basicity increases. B2O3 > Al2O3 > Ga2O3 > In2O3 > Tl2O3 B(OH)3 > Al(OH)3 > Ga(OH)3 > In(OH)3 > Tl(OH)3 Trend in acidity

Basic Character increases Oxides and hydroxide of Al and Ga shows amphoteric nature.

2Al(OH)3(s) + 3H2SO4(aq) Al2 (SO4)3(aq) + 6H2O

Al(OH)3(s) + NaOH(aq) Na[Al(OH)4](aq)

Preparation of oxides of B and Al

(i)

B2O3 is a good dehydrating agent and reacts with water to form orthoboric acid (H3BO3).B2O3 is an acidic oxide dissolves in alkalies to give borates as well as metaborates.

(ii)

It is amphoteric in nature. Crystalline alumina, Al2O3, exists in many forms, one of these called corundum, is very hard and is used as an abrasive.

3. Halides

Boron and aluminium combine with halogens to form trihalides having the general formula MX3.

Halides of boron are covalent, boron trihalides exist as monomer having planar triangular geometry. BX3 acts as lewis acid. Among boron trihalides the order for strong lewis acid is

BF3 < BCl3 < BBr3

Illustration 31. Aluminium has low density than boron why?

Solution: When we move down from B to Al 3rd shell comes which is having 3s, 3p and 3d sub shell but electrons are only present in 3s2 and 3p2 so due to much increase in size as compare to mass its density is lower than boron.

Illustration 32. Why Ga has lowest melting & boiling point exceptionally?

Solution: Gallium has very unusual crystal structure which leads it to have diatomic molecular structure rather having metallic character.

Illustration 33. Why Ga has small size than Al exceptionally?

Solution: When we move down to Ga from Al we see that 3d shell fills before Ga and due to increase in effective nuclear change size of Ga in smaller than aluminium.

Illustration 34. Write the structure for borazine & what is its common name?

Solution:

Its common name is inorganic Benzene

Exercise 31.(i) Why down the group +1 oxidation state becomes more stable in boron family? (ii) What do you understand by banana bond? (iii) Why BF3 acts as Lewis acid? (iv) Why BCl3 is more acidic than BF3 although the electro negativity is high for F

than Cl?

Aluminium

Aluminium although is a reactive metal according to the electrochemical series, it is revered unreactive due to the formation of a oxide film on its surface.

Occurrence

It does not occur in native form. The important ores of aluminium are(i) Bauxite (Al2O3.2H2O) (ii) Cryolite (Na3AlF6) (iii) Feldspar (KAlSi3O8) (iv) Mica KAlSiO10(OH)2

Extraction

Aluminium is extracted from bauxite in a two stage process.

(i) Stage – I It involves the extraction of alumina (Al2O3) from bauxite.

(ii) Stage – II Involves extraction of pure aluminium from Al2O3 by its electrolysis in molten cryolite [Na3AlF6].

(a) Purification of Bauxite

Bauxite contains SiO2, iron oxide and titanium (IV) oxide as impurities. The bauxite ore is digested with a hot concentrated solution of NaOH at about 473 – 523 K and 35 – 36 bar pressure. Aluminium oxide and silica dissolves to form sodium aluminate and sodium silicate respectively leaving behind iron oxide and TiO2 which are filtered off.

The filtrate containing sodium aluminate and sodium silicate is diluted and seeded with freshly precipitated aluminium hydroxide which induces the precipitation of aluminium hydroxide leaving behind sodium silicate in solution.

The aluminium hydroxide is filled dried and calcinated at 1473 K to yield pure alumina (Al2O3)

(b) Electrolysis of pure alumina (Hall – Heroult process) Alumina is obtained from alumina by electrolysis. Pure alumina is dissolved in molten cryolite (Na3AlF6) and is electrolysed in an iron tank lined by carbon internally. Carbon lining serves as cathode, while the number of dipping carbon rods in the fused electrolyte serves as anode. The electrolyte composition ranges are Na3AlF6(80 - 85%), CaF2(5 – 7%), AlF3(5 – 7%) Al2O3

(2 – 8%).

Cryolite improves the electrical conductivity of the cell as alumina is a poor conductor. Also, cryolite lowers the melting point of the mixture to about 1250K. When electric supply is given aluminium is liberated at cathode and gets collected at the bottom of the tank from where it is removed through outlet. Oxygen liberates at anode and combines with the carbon of the anode to produce carbon monoxide which either burns or escape out. At cathode: Al3+ (melt) +3e Al

At anode: O2 O + 2e

C(s) + O CO(g) CO + O CO (g) CO + O CO2 (g)

Properties of Aluminium Physical propertiesAluminium is a soft light silvery – white metal which soon looses its shine due to formation of a layer of oxide on it. It is malleable and ductile. It m.p. is 932K. Electrical conductivity is twice than copper on weight to weight basis. Chemical properties 1. Action of waterIt does not react with pure water but it is corroded by water containing salts. 2. Action of airIt is not affected by dry air but highly electropositive and redily reacts with moist air and forms a hard protective layer of Al2O3, which renders it positive

3. Action of AcidsAluminium dissolves in dilute mineral acids and produces hydrogen.

4. Reaction with strong alkali Aluminium gives aluminate with liberating hydrogen too.

Uses of Aluminium

Aluminium is extensively used in industries as well as in every day life. 1. It is used in making many alloys, utensils, construction aerospace, window angles. 2. It is a good conductor of electricity, thicker cables of aluminium are used for transmission of

electricity. 3. Aluminium foil is used in wrapping cigarettes, confectionary etc. 4. Aluminium is also used in the aluminothermic process for production of chromium and

manganese from their ores.

Illustration 35. Name two alloys of aluminium with their constituents.

Solution: Magnalium (Al and Mg), Duralumin (Al, Cu, Mg and Mn)

Exercise 32.How corundum is obtained from alumina?

Carbon Family The carbon family or group 14 consists of carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn) and Lead (Pb). Carbon and silicon are non metals; germanium is a semi – metal or metalloid whereas tin and lead are metals. All the elements of group 14 have four electrons in their valence shells. Their general electronic configuration is ns2np2. Carbon and silicon are the most abundant elements in the group in earth crust. General introduction to the elements of carbon familyCarbonCarbon is the sixth most abundant element in the known universe but not nearly as common on the earth, despite the fact the living organisms contain significant amounts of the elements. Carbon commonly occurs in environment as methane (CH4) and carbon dioxide (CO2). Carbon exists in several forms called allotropes. Diamond is one of very strong crystal lattice, known as a precious gem. Graphite is another allotrope in which carbon atoms are arranged in planes which are loosely attracted to one another (hence used as lubricant). The recently discovered fullerenes are yet other form of carbon. Carbon has a very high melting and boiling point and rapidly combines with oxygen at elevated temperature. An important (but rare) radioactive isotope of

carbons, C – 14 is used to date ancient objects of organic origin .Silicon The silicon name is taken from Latin silver which means “flint”. The element is on second position in abundance in the earth’s crust after oxygen, was discovered by Berzelius in 1824.

The most common compound of silicon, SiO2 is the most abundant chemical compound in the earth’s crust. Silicon is a crystalline semi – metal or metalloid. One of its forms is shiny, grey and very brittle. In another allotropic form silicon is a brown amorphous powder most familiar in “dirty” beach sand. Germanium (Ge) Like silicon, germanium is used in the manufacture of semi – conductor devices. But unlike silicon, it is rather rare. It was also predicted by Mendeleev in 1871 (ekasilicon) to fill out his periodic table and was discovered in 1886 by Winkler. It is generally extracted from the by products of zinc – refining.

Tin (Sn) It is one of the major elements along with copper used in bronze. It was named after the Etruscan god Tinia, the chemical symbols for it is taken from Latin Stannum . The metal is silvery white and very soft when pure. It has the look of freshly cut aluminium but the feel of lead. Polished tin is slightly bluish. It has been used for many years in the coating of steel cans for food because it is more resistant to corrosion than iron. It is chiefly used in solders. SnF2 is found in fluoride toothpastes. Lead (Pb) Its symbol came from ‘plumber’ word because since old plumbing was done with lead pipes. Although lead is not very common in earth’s crust, what is there is readily available and easy to refine. Its chief use today is in lead – acid storage batteries such as those used in automobiles. In pure form it is too soft to be used for much else. Lead has a blue – white colour when first cut but quickly dulls on exposure to air forming Pb2O. Various isotopes of lead come at the end of the natural decay series of elements like uranium, thorium and actinium. These are Pb – 206, Pb – 207 and Pb – 208.

Illustration 36. Why elemental silicon does not form graphite like structure as carbon does? Explain

Solution: This is due to reluctance of silicon to form pp pp multiple bonds because of large size of silicon atom. Hence silicon exists only in the diamond structure.

Illustration37. Why graphite is used as lubricant?

Solution: Its layers are attached with the weak Vander waal’s forces so that it can slide over each other hence used as lubricant.

Exercise 33.Write the name of two ores of lead.

Exercise 34.Why diamond shines very much?

Exercise 35.Why the other elements except carbon do not show pp pp multiple bonding?

General Trends in Physical properties Sr. No.

Property Carbon Silicon Germanium

Tin Lead

1. Configuration [He] 2s22p2 [Ne]3s23p2 [Ar]4s24p2 [Kr]5s25p2 [Xe]6s26p2

2. Common oxidation state

+4 +4 +4 +4, +2 +4, +2

3. Atomic radius (pm) 77 117 122 140 1754. First ionization

energy (KJ/mol)1086 786 762 709 716

5. Electronegativity 2.5 1.8 1.8 1.8 1.9Catenation

A remarkable property of carbon is its ability to form compounds in which carbon atoms are linked to one another in chains or rings. This property of forming chains and rings is knows as catenation. On going down the group, tendency of catenation decreases. C >> Si > Ge » Sn >> Pb Due to high tendency of catenation carbon forms bonds with other carbon atoms and forms so many compounds which are studied in organic chemistry. Allotropy A characteristic property of the elements of carbon family is that these show allotropy. Example: carbon has two important allotropic forms i.e. diamond and graphite. Allotropy is the existence of an element in two or more forms, which are significantly different in physical properties but have similar chemical properties. In diamond carbon is sp3 hybridised and has four tetrahedral bonds with adjacent carbon atoms. In graphite carbon is sp2 hybridised. The three cavalent bonds form hexagonal layers and fourth unhybridised p – electron of each carbon forms an extended delocalized p - bonding with carbon atoms of adjacent layers. Due to this free electron graphite is electric conductor while diamond is not. Also due to sliding property of graphite it has been using as a lubricant. Chemical Properties of carbon family Due to high ionization enthalpies M4+ ions of the group are not known hence the group elements mostly forms covalent compounds with a covalency of four. Carbon differs distinctly from other elements of group 14 on the basis of its ability to form pp - ppmultiple bonds to itself and to other elements like nitrogen and oxygen. The chemical characteristics of the group 14 elements are discussed below:1. Hydrides Carbon forms large number of cyclic and acyclic hydrides known as hydrocarbons. These have general formulas of type CnH2n+2 (alkanes), CnH2n(alkenes), C2H2n2 (alkynes) etc. Silicon and germanium also forms few hydrides of formula MnH2n+2 (where M = Si, n = 1 to 8 or M = Ge, n = 1 to 5). These are known as Silanes and Germanes respectively. Only one hydride of tin, namely monostannane (SnH4) is known. Thermal stability of these hydrides decreases down the group. The reducing characters of these hydrides increase on going down the group. 2. OxidesThe group forms two types of oxides, monoxides and dioxides. (i) MonoxidesAll except Si forms monoxides of general formula MO, such as CO, GeO, SnO and PbO. CO is only neutral otherwise all are basic in nature. Structure of COThe molecular orbital configuration of CO is , which suggests the

presence of triple bond between carbon and oxygen atoms. However, CO molecule is considered to be a resonance hybrid of the following structures:

CO, is colourless toxic gas and it forms a number of co – ordination compounds with transition metals e.g, Ni(CO)4, Fe(CO)5 and Cr(CO)6. Such compounds are called as organometallic compounds because carbon is bonded to metals. (ii) DioxidesAll the elements of groups 14 forms dioxides of general formula MO2, such as CO2, such as

CO2, SiO2, GeO2, SnO2 and PbO2. (a) Among these oxides CO2 and SiO2 are acidic whereas GeO2, SnO2 are amphoteric in nature

and PbO2 behaves as base.

SnO2, as it is amphoteric reacts with acid as well as base.

(b) PbO2 powerful oxidizing. It readily dissolves in acids forming Pb(II) salts with the liberation of oxygen.

(c) CO2 differs very much from other dioxides. CO2 is gas at room temperature. SiO2 is solid at room temperature. CO2 is monomeric, nonpolar and linear molecule while SiO2 exists in 3-D structure.

(iii) Carbon suboxide (C3O2) In addition to CO an CO2, C3O2 in also formed by carbon.

It is a foul smelling gas with b.p. 6C. It has reaction with H2O or HCl.3. Halides There forms MX4 type halides and except carbon others also forms MX2 type. The stability of MX2, dihalide increases on moving down the group. CX2 << SiX2 << GeX2 << SnX2 << PbX2

Except CX4 all other tetrahalides can be hydrolysed, due to presence of vacant d – orbitals. Due to this reason only tetrahalides of all the elements in group 14 except carbon can react with halogen acids. e.g.

Ionic character and thermal stability decreases with the increase in size of halide ion.

Tin(II) chloride is obtained by dissolving tin in conc. HCl; when the solution is cooled, crystals of tin (II) chloride dehydrate SnCl2.2H2O separate out. Anhydrous SnCl2 is prepared by heating tin in a current of HCl vapour. SnCl2 is used as a reducing agent in acid solution.

Lead (II) halides are formed by adding halide ions to a soluble lead salt.

Pb(II) halides are colourless solids excepts PbI2, which is yellow. They are sparingly soluble in water. The formation of PbCl2 or PbI2 serves as a test for the detection of Pb2+ in qualitative analysis. Silica (SiO2)

Silicon is unable to form pp pp bond with oxygen atom due to its relatively large size. Thus it satisfies its all four valency with four oxygen atoms and constitutes three - dimensional network. In this structure each oxygen atom is shared by two silicon atoms. Three crystalline modification of SiO2 are quartz, cristobalite and tridymite of which quartz and cristobalite are important. Quartz (rock crystal) is the purest form of silica. It is used in preparation of costly glasses and lenses. It is also used as piezoelectric material (crystal oscillators and transducers). Several amorphous forms of silica such as silica gel and fumed silica are known. Silica gel in made by acidification of sodium silicate and when dehydrates, is extensively used as a drying agent in chromatographic and catalyst support.

Structure of Silica

Artificially silica can be obtained by following methods. (i)(ii)(iii)

Properties of Silica Pure silica is colourless but sand is brownish or yellowish due to presence of impurities of iron oxide.

Silicates and Silicones (A) Silicates This is the general term applied for the solids with silicon – oxygen bonds. Some of the silicate minerals are quartz, asbestos (CaMgSi2O6), feldspar (KAlSi3O8), mica [KAl2(Si3AlO10) (OH)2] and zeolites (Na2Al2SiO8.xH2O).The solids contain silicate ion (SiO4)4- as the basic structural unit. The silicate ion is tetrahedral in structure and when the one or more oxygen atoms between such tetrahedrons, a complex structure arise. The silicates may be classified in to chain silicates, ring silicates, cyclic silicates, sheet silicates, three – dimensional silicates depends on the way in which the (SiO4)4- tetrahedral units are linked together.

(b) Silicones Silicones are polymeric compounds containing repeated R2SiO units. The name is given silicone because their empirical formula is analogous to that of ketones (R2CO). Silicones are chemically inert, water repelling nature, heat resistance and having good electrical insulating properties. They are used as sealants, greases, electrical insulators and for water proofing of fabrics. Commercial silicon polymers are usually methyl derivatives and to a lesser extent phenyl derivatives. They are prepared by the hydrolysis of R2SiCl2(R = Me or Ph).

The starting alkyl substituted chlorosilanes are obtained by direct reaction of RCl with silicon in the presence of metallic copper as a catalyst.

It is interesting to note that hydrolysis of alkyl trichlorosilanes, RSiCl3 gives cross linked polymers instead of chain polymers.

Zeolites

Zeolites are microporous aluminosilicates of general formula and may

be considered as open structure of silica in which aluminium has been substituted in a fraction of

of the tetrahedral sites. The aluminosilicates framework is constituted by negative

charge and exchangeable cations of valence n. the void space which is approximately 50% of the volume is occupied by m molecules of water in the unit cell.Extraction of TinThe chief ore of tin is cassiterite (SnO2) which on reduction with carbon produces tin. This ore contains about 10% of tin. It is extracted in following manner:(i) The ore is crushed and washed with water to remove lighter impurities. (ii) It is roasted than to remove impurities such as arsenic and sulphur as volatile oxides. (iii) The roasted ore is heated with coal in a reverberatory furnace at 1500 K.

The crude tin obtained as above is contaminated with iron, lead and other metals. It is therefore remelted on an inclined surface. Tin has a lower melting point hence melt first and flows down leaving other less fusible metals. Molten tin is finally stirred with green poles of wood in contact of air. In this process any remaining metals impurities are oxidized forming a scum which rises to the surface and is removed. Thus process is called poling. Extraction of LeadThe ores of lead are: (i) Galena, PbS (ii)Cerussite, PbCO3 (iii) Anglesite, PbSO4

The chief ore is galena (PbS). The ore is first concentrated by froth – floatation and then roasted in a limited supply of air to give PbO which is reduced to the metal by heating with coke and limestone (flux) in a blast furnace.

In another way, the mixed sulphides (PbS + ZnS) are roasted to obtain oxides. The mixed oxides are reduced to their respective metals with coke by heating in blast furnace.

Molten lead is trapped from the bottom of the furnace. Zinc vapours which come out from the top of the furnace are condensed.

ANSWER TO EXERCISES

Exercise 1:The three general characteristics are:(i) The compounds of the s-block elements are mainly ionic.(ii) The valency is equal to the group number.(iii) Due to low ionization energy, s-block elements are good reducing agents.

Exercise 2:Sodium is the most abundant alkali metal in the earth’s crust.

Exercise 3:This is due to abnormal increase in the atomic size of potassium.

Exercise 4:The Eo value (reduction potential) depends on the three factors i.e. sublimation, ionization and hydration enthalpies. With the small size of its ion lithium has the lightest hydration enthalpy which accounts for its high negative Eo value and its reducing power.

Exercise 5:Lithium

Exercise 6:Potassium is softer than sodium due to weak metallic bonding because of the large size of K atoms.

Exercise 7:Low ionization enthalpy, strongly electropositive nature, tendency to attain noble gas configuration by the loss of one valence electron makes sodium highly reactive.

Exercise 7:Increase in size and low density make s sodium soft.

Exercise 8:(B)

Exercise 9:(i)(ii)(iii)

Exercise 11:Potassium superoxide hydrolyse to form KOH and H2O2 and O2 gas.

Exercise 12:This is due to increase in IE and increase in the size of the cation that the distance between the M – O bond increases. This results in the greater separation of the cation from the hydroxide ion resulting in the greater concentration of the hydroxide in the solution consequently, the basicity increases from LiOH to CsOH.

Exercise 13:This is to due to small size of Li+ and F– ions, which leads to high lattice energy of LiH, hence less soluble in water

Exercise 14:The less solubility of CsI in water is due smaller hydration energy of its ions.

Exercise 15:(D). Basicity of hydroxides of group-I elements increases from top to bottom. Hence (D) is correct

Exercise 16:(D)

Exercise 17:(i) (ii)(iii)

Exercise 18: (B)

Exercise 19:A : CaSO4

B : CaSO4 2H2O

Exercise 20:Magnesium burns in air with dazzling brilliance due to the formation of MgO and Mg3N2.

Exercise 21:The higher IE is more than compensated by the higher enthalpy of lattice formation.

Exercise 22: (A)

Exercise 23:This is due to small size and high electronegativity of beryllium.

Exercise 24:BeCO3 is the least stable alkaline earth metal carbonate. To slow down its decomposition, we store it in an atmosphere of carbon dioxide.

Exercise 25:This is due to the hydrolysis of the beryllium ion.

Exercise 26:The milk of magnesia is magnesium hydroxide. It is used as an antacid to neutralise excess of acid in the stomach.

Exercise 27:(i)

(ii)

Exercise 28:Calcium (or CaO) is used for drying alcohol and non-acidic gases.

Exercise 29:Anhydrous CaCl2

Exercise 30:Mg(OH)2 and MgCO3 are the constituents of toothpaste.

Exercise 31:(i) +1 oxidation state is exhibited when the s – electrons of the valence shall do not

participate in bond formation. This is known as inert pair effect. The inert pair effect becomes more predominant down the group the nuclear charge increases much more than the corresponding increase in distance of outer electrons from the nucleus.

(ii) Banana bond is formed in hydrides of boron (B2H6). In which two B H B bonds are formed called as 3c 2e (three centered 2 electron bond)

(iii) In BF3 boron is having sp2 hybrid state and the three B F bonds are formed by axial overlaps of sp2 hybrid orbital of boron and p – orbital of halogen. Since there are only six electron in the valence shell of boron atom in BF 3, therefore, it has great tendency to accept two more electorns in order to acquire stable octet.

(iv) This is explained on the basis of halogen boron back p - bonding. One of the 2p orbital of F atom having lone pair overlaps side wise with the empty 2p – orbital of boron atom to form pp - pp back bonding. This is also known as back donation. As F has maximum tendency for back donation, in BF3 electorn deficiency of boron gets compensated and its lewis acid character decreases.

Exercise 32:Corundum is obtained by heating amorphous alumina to about 200 K.

Exercise 33:Anglesite (PbSO4) and Cerrusite (PbCO3)

Exercise 34:Due to total internal reflection of light.

Exercise 35:Due to their relatively larger atomic size.

Miscellaneous Exercises

Exercise 1: Complete and balance the following reactions:(i) K + O2

(ii) K + H2O(iii) K + H2

Exercise 2: Write the formulae of(i) Albite (ii) Chile salt petre(iii) Glauber’s salt (iv) Borax

Exercise 3: What happens when NaOH is added to Zn (+II) solution?

Exercise 4: When an alkali metal dissolves in liquid ammonia the solution acquires different colours. Explain the reasons for this type of colour change.

Exercise 5: Name one reagent or one operation to distinguish between(i) BeSO4 and BaSO4 (ii) Be(OH)2 and Ba(OH)2

Exercise 6: Write the chemical formulae of the following minerals.(i) Limestone (ii) Fluorspar(iii) hydroxyapetite (iv) Anhydrite

Exercise 7: What is the mixture of CaCN2 and carbon known as?

Exercise 8: What is ‘potash magnesis’?

Exercise 9: Sodium peroxide forms a white compound when it comes into contact with moist air. Explain.

Exercise 10: How can sodium chloride be purified?

ANSWERS TO MISCELLANEOUS EXERCISES

Exercise 1: (i)(ii)(iii)

Exercise 2: (i) (ii)(iii) (iv)

Exercise 3: First addition of NaOH to Zn(+II) solution gives a white precipitate of Zn(OH)2

which dissolves in excess of sodium hydroxide.

Exercise 4: The solution of alkali metal in liquid ammonia is coloured and metastable.The dilute solutions are blue and the blue colour is due to the presence of solvated electrons.

The blue solution is also paramagnetic in nature:In concentrated solutions, the ammoniated (solvated) metal ions are bound by the free unpaired electron which have been described as expanded metals. The colour is bronze and the solution is diamagnetic.

Exercise 5: (i) BeSO4 is soluble in water while BaSO4 is not.(ii) Be(OH)4 dissolves in NaOH while Ba(OH)2 is insoluble.

Exercise 6: (i) Limestone – CaCO3 (ii) Fluorspar – CaF2

(iii) Hydroxyapetite – Ca5(PO4)3.OH (iv) Anhydrite – CaSO4

Exercise 7: A mixture of calcium cyanamide (CaCN2) and carbon is known as nitrolim. It is used as a fertilizer.

Exercise 8: ‘Potash magnesis’ is a double salt used as a fertilizer. Its formula is K2SO4.MgSO4.6H2O

Exercise 9: Sodium peroxide when exposed to moist air turns white due to formation of NaOH and Na2CO3.

Exercise 10: The salt obtained by various manufacturing processes is contaminated. Hence, to purify sodium chloride, it is first of all treated with ammonium carbonate to precipitate down the carbonates of calcium and magnesium.

The solution is filtered to remove these precipitates and through the filtrate, hydrochloric acid is passed. As a result, sodium chloride settles down. This is found to be quite pure. Only KCl is present as impurity which is removed by repeated crystallisation.

SOLVED PROBLEMS

Subjective:

Prob 1. is soluble in water whereas is insoluble. Why?

Sol. The lattice energy of is less than the hydration energy whereas the lattice energy of the (because of bivalent charge) is very high so that hydration energy released is not sufficient to break the lattice and remains insoluble.

Prob 2. Sodium fire in the laboratory should not be extinguished by pouring water. Why?

Sol. Sodium reacts violently with water producing gas which also catches fire. As a result, the fire spreads rather than being extinguished. Therefore, water should not be used for extinguished sodium fire. Pyrene should be used.

Prob 3. Why potassium carbonate can not be prepared by solvay process?

Sol. This is due to the reason that potassium bicarbonate formed as an

intermediate (when is passed through ammoniated solution of potassium chloride) is highly soluble in water and can not be separated by filteration.

Prob 4. Why alkali metals are not prepared by the reduction of their oxides?

Sol. Alkali metals having lowest values of ionization enthalpy are strongest reducing agents. Thus they can not be reduced by any other element.

Prob 5. Why table salt get wet in raining season?

Sol. Pure NaCl is not hydroscopic but table salt is impure. NaCl contains impurities of . All of these being hydroscopic absorb moisture from air in rainy

season. As a result table salt gets wet.

Prob 6. When Mg metal is burnt in air; a white powder is left behind as ash. What is the white powder?

Sol. Mg on burning in air reacts with oxygen and nitrogen resulting in the formation of MgO and magnesium nitride.

Prob 7. In aqueous solution Li+ ion has the lowest mobility. Why?

Sol. Li+ ions are highly hydrated in aqueous solution which result in decrease in its mobility.

Prob 8. Explain, why lithium is kept rapped in paraffin wax and not stored in kerosene oil?

Sol. It is because lithium is a light metal and therefore it floats at the surface of kerosene oil. To prevent its exposure to air it is kept wrapped in paraffin wax.

Prob 9. Why cesium can be used in photoelectric cell, while lithium can not be?

Sol. Caesium has the lowest while lithium has the highest ionization energy among all the alkali metals. Hence, caesium can lose electron very easily while lithium cannot.

Prob 10. Which alkali metal ion has the maximum polarizing power and why?

Sol. Li+ ion has the maximum polarizing power among all the alkali metal ions. Thus is due to small size of Li+ ion as result of which it has maximum charge/radius ratio.

Prob 11. How aluminate iron is obtained?

Sol. On addition of more alkali to aluminium hydroxide, it dissolves to form aluminate ion.

Prob 12. Why cryolite is added to electrolyte while refining aluminium?

Sol. To make the electrolyte more conducting and to reduce the melting point of the mixture so that aluminium can melt at lower temperatures.

Prob 13. CCl4 can not be hydrolysed but SiCl4 can be. Why?

Sol. Due to availability of vacant d – orbitals in valence shell of their central atom, they can easily extend their coordination number beyond four but this is not possible in case of carbon due to absence of vacant d – orbitals.

Prob 14. SiCl2 is solid while SiCl4 is liquid at room temperature why?

Sol. Dihalides are more ionic than tetrahalides hence having high m.p.

Prob 15. How PbO2 is prepared?

Sol. PbO2 is prepared by treating trilead tetraoxide Pb3O4 with dilute HNO3.

Prob 16. How SnO is obtained?

Sol. SnO is obtained by heating tin oxalate SnC2O4 SnO + CO + CO2

Prob 17. Why SiO2 is solid while CO2 is a gas?

Sol. This is because SiO2 exists in silicate form where Si forms tetrahedron SiO42- ions

which are regularly attached with each other and form a giant network structure hence remains solid.

Prob 18. Write about the composition of following alloys. (i) Bronze (ii) Solder (iii) Bell metal (iv) Cum metal

Sol. (i) Cu 90% ; Sn 10% (ii) Sn 67% ; Pb 33% (iii) Cu 80% ; Sn 20% (iv) Cu 86% ; Sn 10% ; Zn 4%

Prob 19. What is magnesia cement? Give its composition.

Sol. When a saturated solution of magnesium chloride is mixed with magnesium oxide, it sets to a hard mass. This hard mass is known as magnesia cement. Its composition is MgCl2.5MgO.H2O.

Prob 20. Why is it that hydrated chlorides of Ca, Sr and Ba can be dehydrated by heating but those of Be and Mg suffer hydrolysis?

Sol.

MgCl2.6H2O or BeCl2.4H2O on heating suffer hydrolysis due to the small size of Mg2+ (or Be2+).

Objective:

Prob 1. The oxidation state of the most electronegative element in the products of the reaction, with dil. are

(A) 0 and -1 (B) -1 and -2 (C) -2 and 0 (D) -2 and +1

Sol. (D)

Prob 2. Among , , and . Unpaired electron is present in (A) (B)(C) (D)

Sol. (D)

Prob 3. A solution of sodium metal in liquid ammonia is strongly reducing due to the presence of (A) sodium atoms (B) sodium hydride (C) sodium amide (D) solvated electrons

Sol. (D)

Prob 4. The compound insoluble in acetic acid is (A) calcium oxide (B) calcium carbonate (C) calcium oxalate (D) calcium hydroxide

Sol. (C)

Prob 5. Which of the following on heating do not decompose? (A) (B)(C) (D) none

Sol. (C)

Prob 6. The solubility in water of sulphates down the Be groups is Be>Mg>Ca > Sr > Ba. This is due to (A) increase in melting point (B) high ionization energy (C) higher coordination number (D) all of these

Sol. (C)

Prob 7. Solubilites of carbonates decreases down the magnesium group due to decrease in (A) entropy of solution formation (B) lattice energies of solids (C) hydration energy of cations (D) inter-ionic attraction

Sol. (C)

Prob 8. The paramagnetic species is (A) (B)(C) (D)

Sol. (A)

Prob 9. Which of the following substances can be used for drying gas? (A) calcium carbonate (B) sodium carbonate (C) sodium bicarbonate (D) calcium oxide

Sol. (D)

Prob 10. (potassium superoxide) is used in oxygen cylinders in space and submariners because it (A) absorbs and increases content (B) eliminates moisture(C) absorbs (D) produces ozone

Sol. (A)

Prob 11. AlCl3 forms dimmer in vapour phase but BCl3 dose not because (A) In Al there are vacant d orbitals in which it accommodates lone pair from chlorine

atoms (B) In BCl3 there is back bonding (C) There is hydrogen bonding in between two AlCl3 molecules in vapour phase (D) None of the above

Sol. (A)

Prob 12. Non existence of PbI4 can be explained on the basis of (A) Fajan’s rule (B) Inert pair effect (C) Both (A) and (B) (D) None of the above

Sol. (B)

Prob 13. CO2 and SiO2 differ in their physical and chemical properties due to (A) Absence of vacant d orbital in carbon (B) more catenation tendency of carbon (C) Greater electronegativity of carbon (D) All

Sol. (D)

Prob 14. Boric acid is prepared from borax by the action of(A) (B) NaOH(C) (D)

Sol. (A)

Prob 15. Carbon monoxide acts as a Lewis base because it has(A) A double bond between C and O atoms(B) A triple bond between C and O atoms(C) A lone pair of electrons on the C atom(D) Two lone pairs of electrons on the O atom

Sol. (C)

Fill in the blanks Prob 16. Sea water is a major source of ………………… where it is present as

………………………

Sol. Na, NaCl

Prob 17. An alloy of Na with Hg is called …………………………….

Sol. Sodium amalgam

Prob 18. …………………….. is used to neutralise acidity of stomach.

Sol. NaHCO3

Prob 19. In the Castner-Kellner cell, mercury acts as ……………………… in the central and as ………………. in the outer compartments.

Sol. Anode, cathode

Prob 20. Ca2Mg2Si8O22(OH)2 is known as …………………….

Sol. Asbestos

True/False Prob 21. Lithium is the hardest alkali metal.

Sol. True

Prob 22. Alkali metals are strong reducing agents and hence can’t be isolated by reduction of their oxides or compounds.

Sol. True

Prob 23. Sodium reacts vigorously with oxygen to form sodium superoxide.

Sol. False

Prob 24. Pure NaCl isn’t hydroscopic but in the presence of impurities like CaCl2.2H2O and MgCl2.2H2O becomes slightly hygroscopic.

Sol. True

Prob 25. K2SO4.MgSO4.6H2O is a double salt and is known as potash alum.

Sol. False

ASSIGNMENT PROBLEM

Subjective:

Level - O

1. Why alkali metals are normally kept in kerosene oil?

2. Why alkali metals impart colour to the flame?

3. Explain why Lithium chloride has more covalent character than potassium chloride.

4. Why BeCO3 is kept only in an atmosphere of CO2?

5. Why Be and Mg do not give characteristics colour to the flame whereas other alkaline earth metals do give?

6. What is the name of recently developed carbon, i.e., molecule?

7. The name drikold is associated with which chemical species.

8. is less stable than why?

9. Carbon shows properties quite different from other members. What is the cause for it?

10. Why Boron does not forms

11. Lithium has highest ionization energy in group I element, yet it is strongest reducing agent. Why?

12. What happens when?(i) NaH reacts with water. (ii) Sodium reacts with excess oxygen. (iii) Water is dropped over sodium peroxide. Write balanced chemical equation for each.

13. What are boranes? Give the chemical equation for preparation of diboranes.

14. What is carbogen? What is its use?

15. Explain why:(i) is gas while is a solid?(ii) is solid whereas is a liquid?

16. Why are alkali metals strong reducing agents?

17. Sodium can not be obtained by the electrolysis of an aqueous solution of sodium chloride. Explain.

18. Account for the following:(i) Be and Mg do not impart any colour to the flame.(ii) Be has less negative value of the reduction potential (Eo).

19. ‘The chemistry of beryllium is not essentially ionic’. Justify the statement by making a reference to the nature of oxide, chloride and fluoride of beryllium.

20. Why is it that BeSO4 and MgSO4 are soluble but BaSO4 is insoluble in water?

Level – I

1. A white solid is either or . A piece of litmus paper turns white when it is dipped into a freshly made aqueous solution of the white solid. Identify the substance and explain with balanced equation.

2. Why is calcium preferred over sodium to remove last traces of moisture from alcohol?

3. Chlorination of calcium hydroxide produces bleaching powder. Write its chemical equation.

4. What is the difference between milk of lime and lime water?

5. Give reason for decreasing order of conductivity of following. Conductivity of Li+ < Na+ < K+ < Rb+ < Cs+

6. Explain the stability of oxides of alkali metals.

7. Why alkaline earth metals forms divalent ions inspite of the fact that IE2 > IE1 of alkaline earth metals?

8. Li is extracted by electrolysis of a fused mixture of LiCl & KCl. Why KCl is added?

9. State any reason for alkaline earth metals having a greater tendency to form complex than alkali metals.

10. What happens when KO2 reacts with water? Write the balanced chemical equation for the reaction.

11. Write balanced equations for the reactions of the following compound with water:

12. Why superoxides of alkali metals are paramagnetic while normal oxides are diamagnetic?

13. when reacted with water gives off but HCl is not obtained from on reaction with water at room temperature.

14. The crystalline salts of alkaline earth metals contain more water of crystallization than the corresponding alkali metals salts why?

15. Explain why halides of Be dissolve in organic solvents while those of Ba do not.

Level – II

1. What happens when(i) magnesium is burnt in air.(ii) quicklime is heated with silica.(iii) chlorine reacts with slaked lime.(iv) calcium nitrate is heated?

2. Explain why(i) Lithium on being heated in air mainly forms the monoxide and not peroxide. (ii) An aqueous solution of sodium carbonate gives alkaline tests.

3. How is anhydrous magnesium chloride prepared from magnesium chloride hexahydrate?

4. BaO2 is a peroxide but PbO2 is not a peroxide why?

5. Alkaline earth metal impart a characteristic colour to the flame? Explain.

6. Explain (a) The hydroxides of alkaline earth metals are weaker bases than their corresponding alkali

metals. (b) Ca, Br, Sr are stored in paraffin but Be & Mg are not?

7. (a) What is nitrolim and how it is useful? (b) Blue colour of the solution of alkali metals in liquid ammonia fades on standing why?

8. Although boric acid B(OH)2, contains three hydroxyl groups yet it behaves as a mono basic acid. Explain why?

9. Which oxide of carbon is an anhydride of carbonic acid?

10. Carbon monoxide is readily absorbed by ammoniacal cuprous chloride solution but SO2 is not. Explain.

11. Which one is more soluble in diethyl ether, anhydrous or hydrous ? Explain in terms of bonding?

12. Calcium burns in nitrogen to produce a white powder which dissolves in sufficient water to produce a gas (A) and on alkaline solution. The solution on exposure to air produce a solid layer of (B) on the surface. Identify the compounds A & B.

13. Explain why halides of beryllium fume in moist air but other alkaline earth metals halides do not.

14. A piece of burning magnesium ribbon continuous to burn in sulphur dioxide. Explain.

15. Starting from SiCl4, prepare in not more than 4 steps, a linear silicone containing methyl groups only.

Objective:

Level – I

1. The solution of alkali metals as in liquid ammonia is used for reduction of ethylenic double bondd, acetylenic triple bonds to double bonds and aromatic compounds under the name Birch reduction(A) Li (B) Na (C) K (D) all of them

2. Which of the following carbide is known by the name hydrolith? (A) (B)(C) (D)

3. Sodium reacts with water more vigorously than lithium because it has (A) has higher atomic weight (B) is a metal (C) is more electropositive (D) more electronegative

4. The hydration energy of is greater then

(A) (B)

(C) (D)

5. Property of alkaline earth metals that increase with this atomic number is (A) Ionization energy (B) solubility of their hydroxides (C) solubility of their sulphates (D) electronegativity

6. Which of the following oxide is neutral?(A) (B) CO(C) ZnO (D)

7. The correct order of solubility of sulphates of alkaline earth metals are (A)(B)(C)(D)

8. Which of the following metals do not impart colour to the flame? (A) Na (B) Mg (C) Ca (D) Sr

9. A fusion mixture is a mixture of (A) (B)(C) (D)

10. Which of the following alkaline earth metals sulphates is least soluble in water? (A) BaSO4 (B) MgSO4 (C) SrSO4 (D) CaSO4

11. The drying agent – which absorbs carbon dioxide and reacts violently with water is (A) sodium carbonate (B) alcohol (C) conc. H2SO4 (D) calcium oxide

12. The compounds of alkaline earth metals have the following magnetic nature(A) diamagnetic (B) paramagnetic (C) ferromagnetic (D) anti ferromagnetic

13. Which of the following alkali metals halides has the lowest lattice energy? (A) LiF (B) NaCl (C) KBr (D) CsI

14. Which of the following alkali metals ions has the lowest mobility in aqueous solution? (A) Li+ (B) Na+ (C) K+ (D) CS+

15. Which has minimum hydration energy? (A) K+ (B) Li+ (C) Ca+2 (D) Al+3

16. Alcohol dissolves in (A) KCl (B) NaCl(C) RbCl (D) LiCl

17. LiNO3 on heating gives (A) O2 (B) NO2 (C) O2 + NO2 (D) none of the above

18. A compound of sodium which when heated gives CO2 is (A) Na2CO3.10H2O (B) NaHCO3

(C) Na2CO3.7H2O (D) Na2CO3.H2O

19. Which of the following carbonate decompose on heating to evolve CO2? (A) Na2CO3 (B) Li2CO3 (C) K2CO3 (D) Rb2CO3

20. Which of the following has the highest solubility product? (A) KOH (B) CsOH (C) LiOH (D) RbOH

Level – II

1. Which one is used as an air purifier in space craft? (A) quick lime (B) slaked lime (C) potassium superoxide (D) anhydrous CaCl2

2. An aqueous solution of sodium sulphate is electrolysed using inert electrodes. The products at the cathode and anode are respectively(A) H2, O2 (B) O2, H2 (C) O2, Na (D) O2, SO2

3. Carbon dioxide is isosturctral with(A) (B)(C) (D)

4. The correct order of increasing C – O bond length of CO, is (A) (B)(C) (D)

5. On heating sodium metals in a current of dry ammonia, the compound formed is (A) sodium amide (B) sodium azide (C) sodium nitride (D) sodium hydride

6. Anhydrous mixture of KF and HF contains which type of ions(A) K+, H+, F (B) (KF+) (HF) (C) KH+, F (D) K+, HF2

7. In the borax bead test of CO2+, the blue colour of bead is due to the formation of (A) B2O3 (B) Co3B2 (C) Co(BO2)2 (D) CoO

8. When K2O is added to water, the solution is basic because it contains a significant concentration of (A) (B)

(C) (D)

9. Borax structure contains (A) two groups and two groups(B) four groups only (C) four groups only (D) three and one groups

10. Which of the following is used in producing neutrons? (A) Rb (B) Ba (C) Cr (D) Be

11. Reduction of BaSO4 with carbon gives(A) BaSO3 (B) BaS (C) BaS2O3 (D) Ba

12. Nitrolim which is used as fertilizer has the composition(A) CaCN2 (B) CaCN2 + C (C) Ca3N2 (D) Ca(CN)2

13. Magnesium keeps on burning in (A) N2 (B) CO2 (C) NO2 (D) N2 as well as CO2

14. Which of the following are arranged in increasing order of solubilities? (A) CaCO3 < KHCO3 < NaHCO3 (B) NaCO3 < KHCO3 < CaHCO3

(C) KHCO3 < NaCO3 < CaHCO3 (D)CaCO3 < NaHCO3 < KHCO3

15. Which of the following metal ions has more polarizing power?(A) Na+ (B) Ca2+ (C) K+ (D) Be2+

16. The thermal stability of alkaline earth metal carbonates decreases in the order (A) (B)

(C) (D)

17. The stability of the following alkali metal chlorides follows the order(A) (B)(C) (D)

18. Identify the correct order of acidic strengths of , CuO, CaO, (A) (B)(C) (D)

19. The correct order of ionic character is (A) (B)(C) (D)

20. The basic character of the oxides, MgO, SrO, K2O, NiO, CS2O increases in the order(A) (B)(C) (D)

ANSWERS TO ASSIGNMENT PROBLEMS

Subjective:

Level-O

1. Alkali metals, when exposed to atmosphere react with oxygen, moisture and CO 2 present in the air forming oxides, hydroxides and carbonates. In order to prevent the reaction the alkali metals are stored in kerosene oil.

2. Alkali metals have low I.E. Their valence electrons absorbs energy from the flame and gets excited to higher energy level When they return to the ground state, the energy is emitted back in the form of radiations frequencies of which fall in the visible range imparting a characteristic colours to flame.

3. Lithium ion due to its small size, has high polarizing power. Thus it can polarize the electron cloud of chloride ion to much greater extent as compared to potassium ion. Hence LiCl has more covalent character.

4. The temperature of decomposition of BeCO3 is less than 373K. Hence it is unstable and can be kept only in an atmosphere of CO2.

5. Be and Mg have high ionization energies. Their valence electrons are not easily excited.

6. Fullerenes. C60 is also known as Buck minister fullerene.

7. Solid carbon dioxide is called dry ice and has drikold as commercial name.

8. This is because Pb is more stable in +2 state due to inert pair effect while Sn is better stable in + 4 state as compared to Pb+4.

9. This is due to small size, absence of d – orbital and property of catenation.

10. This is due to the fact that the sum of all three ionization energies for Boron is quite high.

11. The strength of an element as reducing agent depends not only on ionization energy but also on heat of hydration and heat of sublimation of the element. Li+ ion has very high hydration energy due t its small size which more than compensates for the higher value of I.E. of lithium. As a result is Li behaves as strongest reducing agent.

12. (i) NaH reacts with water to form sodium hydroxide and hydrogen gas.

(ii) Sodium on reaction with excess oxygen forms sodium peroxide

(iii) Na2O2 on reaction with water is hydrolysed to sodium hydroxide and hydrogen peroxide.

13. The hydrides of boron are known as boranes

14. It is a mixture of oxygen and carbon dioxide. It is given to pneumonia patients and victims of the carbon monoxide poisoning.

15. (i) Carbon atom being small in size forms double bonds with O – atoms (pp pp bonding) and hence exists as monomeric linear molecules, silicon atom being large in size form single bonds and is linked to four O – atoms forming a three – dimensional network structure.

(ii) Dihalides are more ionic than tetrahalides.

16. The ionization enthalpies of the alkali metals is very low. It is because of this reason that they form unipositive ions very easily. This implies that they have a great tendency to donate an electron to another atom e.g.

The alkali metals have low value of Eo. The Eo value depends on the three parameters i.e. sublimation, ionization and hydration enthalpies.

or

17. Sodium metal cannot be isolated by the electrolysis of their aqueous solutions because the metal deposited at the cathode, reacts at once with water to form metal hydroxide and hydrogen gas. Thus, at the cathode, metals are not liberated but only hydrogen gas is evolved.

18. (i) Be and Mg do not impart any colour to the flame because the electrons are very strongly bound to the nucleus and very high amount of energy is required to excite the electrons to higher energy level. When these electrons drop back to the original state, energy emitted falls outside the visible region. Hence, no colour to the flame.

(ii) The less negative value for Be arises from the large hydration energy associated with the small size of Be2+ and the relatively large value of the enthalpy of atomization of the metal.

19. Due to exceptionally small size and high ionization enthalpy, the compounds of beryllium are largely covalent. For example:BeO is amphoteric while oxide of other group 2 elements are basic in nature. BeO dissolves in alkali to form beryllate

BeCl2 is largely covalent and is soluble in organic solvents.BeF2 is also highly covalent and due to high lattice enthalpy, the solubility is very less.

20. The greater hydration energies of Be2+ and Mg2+ ion overcome the lattice energy of the salts and their sulphates are soluble. In BaSO4, lattice energy is higher than the hydration energy and hence it is insoluble.

Level – I

2. Both Na and Ca react with water forming their respective hydroxide. In contrast, Na reacts with alcohol to form sodium alkoxide but Ca does not.

3. Bleaching powder is obtained by passing Cl2 into Ca(OH)2. Though bleaching powder is often written as Ca(OCl)2. It is actually a mixture of

4. A suspension of slaked lime, i.e Ca(OH)2 in water is called milk of lime but a clear decanted solution of slaked lime in water is called lime water.

5. Ions are hydrated. Since Li is very small it is heavily hydrated. This makes the radius of the hydrated ions large and hence it move only slowly (although Li+ is very small) and the radius of hydrated Cs+ ion is smaller than the radius of hydrated Li+.

7. This is due to their high lattice energies in the solid state and high hydration energies in aqueous solution.

8. KCl is added to increase the conductivity of LiCl and to lower the fusion temperature.

9. Because of smaller size and higher charge on alkaline earth metals cations as compared to the corresponding alkali metal cations, alkaline earth metals cations have greater tendency to form complexes.

10.

11.

12. Mg burns in air to form MgO and Mg3N2

Magnesium nitride on hydrolysis with H2O give NH3

13. Mg3N2 is a salt of a strong base, Mg(OH)2 and a weak acid (NH3) and hence gets hydrolysed to give NH3. In contract MgCl2 is salt of strong base Mg(OH)2 and strong and HCl and does not undergo hydrolysis to give HCl.

14. Due to smaller size and higher nuclear charge alkaline earth metals have a higher tendency than alkali metals to attract H2O molecules and thus contain more water of crystallization than alkali metals. e.g. LiCl2.H2O & MgCl2.6H2O

15. Halides of Be are covalent because of high I.E. of Be while those of Ba are ionic due to low I.E. of Ba.

Level – II

1. (i)

(ii)(iii)(iv)

2. (i) Li+ ion is smaller in size. It is stabilized more by smaller anion, oxide ion (O 2) as

compared to peroxide .

(ii) An aqueous solution of sodium carbonate gives alkaline tests because Na2CO3

undergoes hydrolysis forming sodium hydroxide.

3. Anhydrous MgCl2 can not be prepared by simply heating MgCl2.6H2O because it get hydrolysed by its own water of crystallization.

However if hydrated magnesium chloride is heated in an atmosphere of HCl gas at 650 K, it checks the above hydrolysis reaction and the hydrated magnesium chloride now loses water of crystallization to form anhydrous MgCl2.

4. Metallic oxides which on treatment with dilute acids produce hydrogen peroxide are called

peroxide. All peroxide contain a peroxide ion having the structure

O O . PbO2 does not contain a peroxide ion and it cannot be called as peroxide.

5. When alkaline earth metals or their compounds are put into a flame, the electron absorbs energy and are excited to higher energy levels. When they returns to their ground state (normal state), they emit the absorbed energy in the form of radiations having particular wavelength.

6. (a) Due to high ionization energies and less solubilities of alkaline earth metals as compared to alkali metals.

(b) Ca, Ba, Sr are much reactive towards oxygen and forms respective oxide but Be & Mg are not so reactive and they form a protective layer of oxide on their surface so they are not stored in paraffin.

7. (a) The mixture of calcium cyanamide and carbon is called nitrolim and is used as show acting nitrogenous fertilizers as it hydrolysis slowly over a period of months evolving NH3

gas.

(b) On standing the ammoniated ions and electrons combine to form metal amide, hence the colour fades.

8. Because of the small size of boron atom and presence of only six electron in its valence shell in B(OH)3, it coordinates with the oxygen atom of the H2O molecule to form a hydrated species.

In this hydrated species B3+ ion because of its small size it has a high polarizing power and hence pulls the electron of the coordinated oxygen atom towards it. The coordinated O – atom, in turn pulls the s - electron of the O – H bond, there by facilitating the release of a proton. As a result, B(OH)3 acts a weak monobasic lewis acid and thus reacts with NaOH solution to form sodium metaborate.

9. Since carbonic acid decomposes to give CO2 and H2O, therefore, CO2 is regarded as an anhydride of carbonic acid.

10. Due to the presence of a lone pair of electrons on carbon in CO, it acts as a lewis bax (or ligand) and thus forms a soluble complex with ammonical cuprous chloride solution.

On the other hand, CO2 does not act as a Lewis bax since it does not have a lone pair of electrons on the carbon atom and hence does not dissolve in ammoniacal cuprous chloride solution.

11. Anhydrous AlCl3 because ethers are non – polar compounds. In hydrous AlCl3 the d – orbitals of Al are already occupied by H2O molecules but in anhydrous AlCl3 d – orbitals of Al are vacant and oxygen of ethers can gives its lone pair to there vacant d – orbital.

12. Ca burns in air to form CaO and Ca3N2. Calcium nitride on hydrolysis with H2O gives ammonia (A).

The alkaline solution of Ca(OH)2 thus formed reacts with CO2 present in the air to form CaCO3 (B)

Thus A = NH3 B = CaCO3 white powder

13. BeCl2 reacts with the moisture to form HCl while other halides do not. BaCl2 + 2H2O Be(OH)2 + 2HCl

14. A piece of Mg ribbon continues to burn in SO2 since it reacts to form MgO and S

This reaction is so much exothermic that heat evolved keeps the Mg ribbon burning.

15.

Objective:

Level - I

1. D 2. B 3. C4. B 5. B 6. B7. B 8. B 9. B10. A 11. D 12. A13. D 14. A 15. A16. D 17. C 18. B19. B 20. B

Level - II

1. C 2. A 3. A&C4. D 5. A 6. D7. C 8. C 9. B10. D 11. B 12. B13. D 14. D 15. D16. A 17. A 18. C19. A 20. C