s-BLOCK ELEMENTS, Exams of Arabic

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CHAPTER

s-BLOCK ELEMENTS

Animation 2.1 : s block elements Source and Credit: eLearn.Punjab

2.1 INTRODUCTION

IN THIS CHAPTER YOU WILL LEARN

The alkali and alkaline-earth metals include the most reactive elec- tropositive elements and a study of their electronic configurations will help in understanding their properties.

The name alkali came from Arabic, which means ‘The Ashes’. The Arabs used this term for these metals because they found that the ashes of plants were com- posed chiefly of sodium and potassium. Alkali metals include the elements, lith- ium, sodium, potassium, rubidium, caesium and francium. These are very re- active metals, produce strong alkaline solutions with water.The alkaline-earth metals are beryllium, magnesium, calcium, strontium, barium and radium.They are called alkaline-earth because they produce alkalies in water and are widely distributed in earth’s crust.

1. To write the electronic configuration of s-block elements in sequence.
2. The occurrence of group IA and IIA elements and the peculiar behaviours of lithium and beryllium.
3. The difference in the physical properties of group IA and IIA elements as well as the differences in the chemical behaviour of their compounds.
4. The commercial preparation of sodium.
5. How sodium hydroxide is commercially prepared.
6. The role of gypsum and lime in agriculture and industry.

The s-block elements are the metals in Group IA and Group IIA of the peri- odic table.They are called the s-block elements because s-orbitals are being filled, in their outer most shells.The elements of group IA except hydrogen are called “Alkali metals” while those of IIA are named “Alkaline-earth metals”.

Alkaline earth metals have two electrons in ‘s’ orbital of their valence shell. All alkaline earth metals lose their two electrons to form dipositive ions M2+, because their ionization energy values are low. They form ionic compounds and show + 2 oxidation state.

The electronic configurations and some physical constants of alkaline earth metals are given in Table 2.2.

Alkaline-Earth Metals

Table 2.2 Electronic Configurations and Physical Constants of Alkaline- Earth Metals

Properties Be Mg Ca Sr Ba Atomic number 4 12 20 38 56 Electronic configurations 1s 2 2s^2 [Ne]3s^2 [Ar]4s^2 [Kr]5s^2 [Xe]6s^2 Ionization energy (kJ/ mol) 899 738 590 549 503 Electron affinity (kJ/mol) 240 230 156 168 52 Electronagetivity 1.5 1.2 1.0 1.0 0. Atomic radius 89 136 174 191 198 Ionic radius of 2+ion (pm) 31 65 99 113 135 Melting points (°C) 1289 649 839 769 725 Boiling points (°C) 2970 1107 1484 1384 1640 Density gm/cm 3 at (20°C) 1.85 1.74 1.55 2.6 3.

Heat of hydration (kJ/mol) 2337 1897 1619 1455 1250

In going down a group the number of shells increase by one at each step and equal to the number of the period to which the element belongs.

Animation 2.3 : s-block elements Source and credit: Crescen

2.1.2 Occurrence of Alkali Metals

Due to high reactivity, the alkali metals occur in nature in the combined state. None of the alkali metals is found free in nature. Sodium and potassium are abundant alkali metals and each constitute about 2.4 percent of earth’s crust. Most of the earth’s crust is composed of insoluble alumino-silicates of alkali metals.

Name of Mineral Chemical Formula

Lithium Spodumene LiAl(SiO 3 ) (^2) Sodium Rock Salt (Halite) NaCl Chile saltpetre NaNO (^3) Natron Na 2 CO 3 .H 2 O Trona Na 2 CO 3 .2NaHCO 3 .2H 2 O Borax Na 2 B 4 O 7 .10H 2 O Potassium Carnallite KCl.MgCl 2 .6H 2 O Sylvite KCl Alunite(Alum Stone)! K 2 SO 4 , Al 2 (SO 4 ) 3 .4Al(OH) 3

Table 2.3 Common Minerals of The Most Important Alkali Metals

Name of Mineral Chemical Formula

Beryllium Beryl Be 3 Al 2 (SiO 3 ) 6 Chrysoberyl Al 2 BeO 4 Magnesium Magnesite MgCO 3 Dolomite MgCO 3. CaCO 3 Carnallite KCl.MgCl 2 .6H 2 O Epsom salt MgSO 4 .7H 2 O Soap stone (talc) H 2 Mg 3 (SiO 3 ) (^4) Asbestos CaMg 3 (SiO 3 ) 4 Calcium Calcite (Lime Stone) CaCO (^3) Gypsum CaSO (^) 4.2H 2 O Fluorite CaF (^2) Phosphorite Ca 3 (PO 4 ) 2 Strontium Strontionite SrCO (^3) Barium Barite BaSO (^4)

Table 2.4 Common Minerals of the Alkaline-Earth Metals

2.1.4 Peculiar Behaviour of Lithium

In many of its properties, lithium is quite different from the other alkali metals.This behaviour is not unusual, because the first member of each main group of the periodic table shows marked deviation from the regular trends of the group as a whole.

The deviation shown by lithium can be explained on the basis of its small radius and high charge density. The nuclear charge of Li+^ ion is screened only by a shell of two electrons. The so-called ‘anomalous’ properties of lithium are due to the fact that lithium is unexpectedly far less electropositive than sodium. Some of the more important differences of lithium from other alkali metals are listed below:

1. Lithium is much harder and lighter than the other alkali metals.
2. The lithium salts of anions with high charge density are generally less soluble in water than those of the other alkali metals, e.g. LiOH, LiF, Li 3 PO 4 , Li 2 CO 3.
3. Lithium forms stable complex compounds, althongh complex formation generally is not a property of alkali metals. One of the stable complexes formed by lithium is [Li(NH 3 ) 4 ]+
4. Lithium reacts very slowly with water, while other alkali metals react violently.
5. Lithium salts of large polarizable anions are less stable than those of other alkali metals. Unlike other alkali metals lithium does not form bicarbonate, tri-iodide or hydrogen sulphide at room temperature.

Animation 2.5 : ALKALI METALS Source and Credit: Docbrown

6. When burnt in air lithium forms only normal oxide, whereas the others form peroxides or superoxides.
7. Lithium hydride is more stable than the hydrides of other alkali metals.
8. Lithium compounds are more covalent, that is why its halides are more soluble in organic solvents and the alkyls and aryls of lithium are more stable than those of other alkali metals.
9. Lithium is the least reactive metal of all the alkali metals.

1. Beryllium metal is almost as hard as iron and hard enough to scratch glass. The other alkaline earth metals are much softer than beryllium but still harder than the alkali metals.
2. The melting and boiling points of beryllium are higher than other alkaline earth metals. (Table 2.2)
3. As reducing agents, the group IIA metals are all powerful enough to reduce water, at least in principle. However, with water, beryllium forms insoluble oxide coating that protects it from further attack.

2.1.5 Peculiar Behaviour of Beryllium

4. Beryllium in particular is quite resistant towards complete oxidation, even by acids, because of its BeO coating.
5. Beryllium is the only member of its group which reacts with alkalies to give hydrogen. The other members do not react with alkalies.

(s) (aq) 2 2 (aq) 2 (g)

Sodium beryllate

Be + 2NaOH → Na BeO + H

2.2 GENERAL BEHAVIOUR OF ALKALI METALS

The reducing property of an element depends on the magnitude of its ionization energy. Reducing agent is a substance which can lose electrons. Since alkali metals have got low ionization energies, so they are strong reducing agents. They are highly electropositive. They react readily with halogens giving alkali metal halides.

Beryllium is the lightest member of the series and differs from the other group IIA elements in many ways.This is due to its small atomic size and comparatively high electronegativity value. The main points of difference are:

1. Low ionization energies make the alkali metals, the most reactive family of metals.
2. Very high second ionization energies indicate that oxidation number higher than 1, are ruled out for the alkali metals.
3. The cations of alkali metals have low charge and large radii than the radius of any cation from the same period, so the lattice energies of their salts are relatively low. Consequently, most of the simple salts of the alkali metals are water-soluble. Most of the salts are dissociated completely in aqueous solution and the hydroxides are among the strongest bases available.
4. They react with oxygen and the surface is tarnished due to the oxides formed. Only lithium burns in air to form the normal oxide, Li 2 O (white solid).

2 2

Lithium oxide

4 Li (s) + O (g) → 2 Li O(s)

2.2.1 Trends in Chemical Properties of Alkali Metals

The exposed metals are oxidized almost immediately by oxygen in air, and in the presence of moisture. The oxides formed react with CO 2 in the atmosphere to form carbonates.

2 2 2 3

Lithium oxide Lithium carbonate

Li O(s) + CO (g) → Li CO (s)

Sodium will undergo a similar reaction, but only if the supply of oxygen is limited. In the presence of excess of oxygen, sodium forms the pale yellow peroxide.

LiH(s) + H O 2 ( )  → LiOH (aq) + H 2 (g)

Due to the presence of hydride ion (H ), the ionic hydrides are used as powerful reducing agents.

7. Lithium is the only Group IA metal that combines with nitrogen and carbon to form nitride and carbide, respectively.

2 3 Lithium nitride

6Li (s) + N (g) → 2Li N(s)

4 Lithium carbide

4Li(s) + C (s) → Li C(s)

Alkali metals react easily with halogens to give halides. Lithium and sodium, for example, react slowly with chlorine at room temperature. Molten sodium burns with a brilliant yellow flame in a chlorine atmosphere to form sodium chloride.

2Na (s) + Cl 2 (g) → 2NaCl(s)

Potassium, rubidium and caesium react vigorously with all the halogens, forming metal halides. All alkali metals form their sulphides when treated with molten sulphur. The general reaction is:

2M (s) + S(s) → M S 2 (s)

2.2.2 Trends in Chemical Properties of Alkaline-Earth Metals

1. The alkaline-earth metals burn in oxygen to form oxides or in the case of barium, the peroxide. Beryllium is the least reactive metal in the group. It is resistant to complete oxidation and stable in air at ordinary temperature but oxidizes rapidly at about 800“C. Therefore beryllium is not tarnished by atmospheric attack but the metal soon loses the silvery appearance.

When exposed to air magnesium quickly becomes coated with the layer of MgO. This layer protects the surface from further corrosion at ordinary temperature.

When magnesium is burnt in air a small amount of nitride is also formed along with magnesium oxide: When barium is heated in air or oxygen at 500 - 600°C, its peroxide is formed.

800 C 2Be(s) + O 2 (g) → 2BeO(s)



2Mg (s) + O 2 (g) → 2MgO(s)

500 600 C 2 2 Barium peroxide

Ba (s) + O (g) − → BaO (s)



2. Hydrides are produced by treating the molten alkaline earth metals with hydrogen, usually under high pressures. Magnesium reacts with hydrogen at high pressure and in the presence of a catalyst (Mgl 2 ) forming magnesium hydride.

Beryllium does not react with water even at red hot temperature but remaining alkaline earth metals produce hydroxides with water.

6. Magnesium is more reactive than beryllium, even though it is not attacked by cold water. Magnesium reacts slowly with boiling water and quite rapidly with steam to liberate hydrogen.

Alkali metal oxides dissolve in water to give strong alkaline solutions. For example:

2.2.3 General Trends in Properties of Compounds of Alkali and Alkaline

Earth metals

i) Oxides

2 2

2 2 2 2

Li O(s) + H O( 2LiOH (aq)

2Na O (s) + 2H O( 4NaOH (aq) + O (g)

l)

l)

100 C 2 2 Steam

Mg (s) + H O(g) → MgO(s) + H (g)



100 C

M (s) + 2H O 2 (l) → M(OH) 2 (s) + H 2 (s)



The reaction of an alkali metal oxide with water is an acid-base reaction and not an oxidation reduction reaction since no element undergoes a change in its oxidation number. The reaction simply involves the decomposition of water molecule by an oxide ion.

O (aq) + H O 2 ( 2OH (aq)

l) →

The basic character of alkali metal oxides increases down the group. Potassium superoxide (K0 2 ) has a very interesting use in breathing equipments for mountaineers and in space craft. It has the ability to absorb carbon dioxide while giving out oxygen at the same time.

The solubility of alkaline earth metal oxides in water increases down the group. BeO and MgO are insoluble but CaO, SrO and BaO are soluble and react with water to form the corresponding hydroxides.

The basic character of the oxides of alkaline earth metals increases down the group. The tendency for group IIA oxides to form alkaline solution is relatively less than that of alkali metals.

4KO (s) + 2CO (g) 2 2 → 2K CO (s) + 3O (g) 2 3 2

Animation 2.6 : Reaction with acids Source and Credit: Learn

BeO is amphoteric in nature since it reacts with both acids and bases.

iii) Carbonates

The carbonates of alkali metals are all soluble in water and are stable towards heat except Li 2 CO 3 which is not only insoluble but also decompose on heating to lithium oxide. The decomposition is made easy because the electrostatic attraction in converting from carbonate to oxide is considerable. In case of large cation like K+^ in K 2 CO 3 , the gain in electrostatic attraction is relatively much less and the decomposition is difficult. Sodium carbonate is very important industrial chemical. At temperature below 35.2°C, Na 2 CO 3 crystallizes out from water as Na 2 CO 3 .10H 2 O, which is called washing soda. Above this temperature it crystallizes as Na 2 CO 3. H 2 O. On standing in air, Na 2 CO 3 .10H 2 O slowly loses water and converted to a white powder Na 2 CO 3 .H 2 0. The solution of Na 2 CO 3 in water is basic due to hydrolysis of carbonate ion.

Unlike the alkali metal carbonates, the alkaline earth metal carbonates are only very slightly soluble in water, with the solublity decreasing down the group. They also decompose on heating and the ease of decomposition decreases down the group.

Na CO 2 3 (s) + 2H O 2 ( )  → 2NaOH(aq) + H CO 2 3 (aq)

CaCO (s) 3 → CaO(s) + CO (g) 2

The ease of decomposition can be related to the size of the metal ion, the smaller the ion, the more is the lattice energy of the resulting oxide and hence higher the stability of the product.

iv) Nitrates

Nitrates of both alkali and alkaline-earth metals are soluble in water. Nitrates of Li, Mg , Ca and Ba decompose on heating to give O 2 , NO 2 and the metallic oxide whereas nitrates of Na and K decompose to give different products.

3 2 2 2

3 2 2 2

3 2 2 2

3 2 2

4LiNO (s) 2Li O (s) + 4NO (s) + O (g)

2 Mg(NO ) (s) 2MgO (s) + 4NO (g) + O (g)

2Ca(NO ) (s) 2CaO (s) + 4NO (g) + O (g)

2NaNO (s) 2NaNO (s) + O (g)

→

→

→

→

v) Sulphates

All the alkali metals give sulphates and they are all soluble in water. The solubilities of sulphates of alkaline earth metals, gradually decrease down the group. BeSO 4 and MgSO 4 are fairly soluble in water. CaSO 4 is slightly soluble, while SrSO 4 and BaSO 4 are almost insoluble.

Calcium sulphate occurs in nature as gypsum CaSO 4 .2H 2 O. When it is heated above 100°C, it loses three quarters of its water of crystallization, giving a white powder called’ Plaster of Paris.

4 2 4 2 2 2

Gypsum Plaster of Paris

2CaSO. 2H O → (CaSO ). H O + 3H O

2.3 COMMERCIAL PREPARATION OF SODIUM BY DOWNS CELL

Most of sodium metal is produced by the electrolysis of fused sodium chloride. Since the melting point of sodium chloride is 801°C, some calcium chloride is added to lower its melting point and to permit the furnace to operate at about 6000 C.