Chemistry Part Ii Chapter 10 The S-Block Elements

Alkali metal has one electron each in the valence subshell of their atoms. Since they have only one electron in valence subshell, therefore, they lose easily, owing to their low ionisation energies. Therefore, alkali metals are highly reactive chemically and do not exist in the free or native state.

It is due to similarity in their:
(i) charge/size ratio.   
(ii) ionic sizes.

These are generally covalent.

It is due to high hydration energy of small lithium ion which more than compensates high ionisation enthalpy. The hydration energy of Li⁺ ion is very high and as a result of this reduction potential of lithium is very low(E°= –3). So lithium acts as the most powerful reducing agent in aqueous solution.

Caesium is the most reactive alkali metal, as it has lowest first ionisation enthalpy and lowest electron negativity. 

It is because after the loss of one electron, alkali metal ion (M⁺ ) is formed which has stable noble gas configuration and it is very difficult to remove another electron to form M²⁺ ion as the second ionisation enthalpy is very high. For example, when lithium loses one electron and forms Li⁺ which is similar to Helium atom.
 

This is because in the air they are easily oxidised to oxides which may dissolve in the moisture of the air to form hydroxides or they also combine directly with water vapours and catch fire. 

Lithium cannot be stored in kerosene oil because it is the lightest metal and it floats on its surface and thus reacts with air.

Lithium form complexes because it is comparatively smaller in size and has a high tendency to accept electrons.

Oxidation state of Oxygen Na₂O₂ is =(-1)
 Let x be the oxidation state of Na in Na₂O₂. Since peroxide linkage is present in Na₂O₂ in which the oxidation of O is –1.

Hence oxidation state of Na in  is +1. 

When magnesium is burnt in air, dazzling brilliant light is given out and a mixture of magnesium oxide and magnesium nitride is formed.

Alkali metals form a blue coloured solution with ammonia due to the formation of ammoniated electron.

Tips: -

Hydrogen gas is liberated.

Sodium reacts with water with vigorous evolution of hydrogen which catches fire. 

Sodium metal is extracted by Down's process

A concentrated solution of sodium chloride in water is called brine. 

Table salt is crude sodium chloride and contains impurities of calcium chloride, magnesium chloride and magnesium sulphate. Since these are deliquescent in nature they absorb moisture during the rainy season from the air. As a result, table salt gets wet.

(i) Baking soda: NaHCO₃
(ii) Salt cake: Na₂SO₄.

On heating, sodium bicarbonate decomposes to evolve CO₂ while sodium carbonate does not.

A mixture of Na₂CO₃ and K₂CO₃ is known as fusion mixture.

Fr(Z = 87);  Ra(Z = 88)

It is due to small atomic size and high value of ionisation enthalpy. 

Beryllium halides on exposure to moist air undergo hydrolysis to form halogen acids. As a result, these halides fume in air.

Alkaline earth metals do not occur in the free state because of their high reactivity.

Mg⁺ ion unstable because Mg⁺ ion has still the tendency to lose another electron to form Mg²⁺ ion which has the stable noble gas configuration of the nearest inert gas (Neon).

A mixture of Na₂O₂ and dilute HCl is called soda bleach, commercially known as oxone.

Sequence of alkaline earth metals in the group:
Be, Mg, Ca, Sr, Ba and Ra.

Radium

Reactivity of an element depends on upon the ionisation enthalpy. Smaller the ionisation enthalpy, greater is the reactivity. As potassium has smaller ionisation enthalpy than calcium. Thus, potassium is more reactive than calcium.

It is due to increase in effective nuclear charge. Alkaline earth metals and their ions are smaller in size than those of group 1.

Alkaline earth metals have less electropositive character than the corresponding alkali metals because of higher ionisation enthalpy values of alkaline earth metals as compared to alkali metals.

The solubility of alkaline earth metal compounds are comparatively less than the corresponding alkali metal due to the small size of bivalent  ions of alkaline earth metals, the lattice enthalpy of alkaline earth metal compound is very high as compared to those of the metals of the first group.

BeCl₂ < MgCl₂ <  CaCl₂ < KCl.

BeCl₂ is a covalent compound.

Hydride of calcium, CaH₂ is known as hydrolith.

First ionisation enthalpy IE₁: Na < Mg
Second ionisation enthalpy : IE₂: Na > > Mg

Be and Mg metals forms complexes.

It is due to its polymeric nature.

Be(Z = 4) which has electronic configuration 1s² 2s² shows electronic configuration of 1s² 2s¹ 2px¹ in excited state and one s-orbital (2s) and one p-orbital (2p) can intermix to form two sp hybrid orbitals which can be used to form BeCl₂.

D.

It is due to:
(i) similar size of atoms or ions
(ii) similar electronegativity and
(iii) similar polarising power.

Diagonal relationship

This is due to the hydrolysis of the beryllium ion.

Be²⁺ +2H₂O --> Be(OH)₂ + 2H⁺

Ba(NO₃)₂.

A solution containing 12 g of MgCO₃ per 100 cc of water containing dissolved CO₂ is known as 'Fluid Magnesia'.

It is anyhydrous calcium sulphate obtained by heating gypsum above 393 K.

A mixture of slated lime, sand and water.

Magnesium.

Gypsum CaSO_4.2H₂O

The cement is a mixture of magnesium oxide (burnt magnesia) with magnesium chloride with the approximate chemical formula Mg₄Cl₂(OH)₆(H₂O)₈, corresponding to a weight ratio of 2.5–3.5 parts MgO to one part MgCl₂. A variant uses zinc oxide with zinc chloride instead of the magnesium compounds.

It continues to burn.

Bones are made of calcium phosphate. 
Therefore, phosphate  ions are the ions associated with them.

When calcium oxide(quick lime) is heated with coke in an electric furnance at 2273- 3273 K, calcium carbide (CaC₂) is formed. 

Both Na and Ca react with water to form their respective hydroxides. In contrast, Na reacts with alcohol to form sodium ethoxide but calcium does not.

Therefore, calcium is preferred over sodium.

It is a mixture of cement, sand, gravel (small pieces of stone) and appropriate amount of water.

It is a waste product from the steel industry and has properties similar to that of cement. It mainly consists of calcium silicate.

The elements of s-block and p-block are collectively called a representative or main group elements. The elements of groups 1 and 2 (s-block), 13 to 18 (p-block) constitute the main group elements or representative elements.

The metallic elements lithium (Li), sodium (Na), potassium(K), rubidium (Rb), caesium (Cs) and francium (Fr) constitute group 1 of the periodic table. Francium is a radioactive element. They are known as alkali metals because their hydroxides are strong alkalies or bases.
Electronic configuration: The atoms of alkali metals have one electron in s-orbital outside a noble gas core. Therefore, their general electronic configuration is:
[Noble gas]ns¹ where n = 2 to 7
The electronic configurations of alkali metals are given below:

The 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, there is a progressive addition of new energy shells. Although, the nuclear charge also increases down the group yet the effect of the addition of new shells is more predominant.
The radii of positive ion are smaller as compared to that of parent atom. Within the group, the ionic radii increase with the increase in atomic number.

Alkali metals possess lowest ionisation enthalpies in their respective periods. However, within the group, the ionisation enthalpies of alkali metals decrease down the group.
Reason: The atoms of alkali metals are largest in their respective periods, therefore, the outermost electrons, which are far away from the nucleus, experience a less force of attraction with the nucleus and hence can be easily removed. Decrease in ionisation enthalpy, on moving down the group, is due to the increase in the size of the atoms of alkali metals and increase in the magnitude of screening effect is by virtue of an increase in the number of intervening electrons.
Electropositive character: On account of their low ionisation enthalpies, these metals have a strong tendency to lose their valence electrons and thus change into positive ions. Consequently, alkali metals are strongly electropositive or metallic in character. As this tendency for losing electron increases down the group, the electropositive character increases.

Alkali metals have metallic bonds in which the positive ions are held together by electrons. The valence electrons move freely from one metal ion to another without much difficulty and thus have high electrical conductivity. Solubility depends on lattice enthalpy and hydration enthalpy. Since the alkali metal ions have low lattice enthalpy and high hydration energy so they form soluble salts.

Due to low values of first ionisation enthalpies, low heats of atomization and high hydration enthalpies, alkali metals lose their valence electron readily and hence are chemically highly reactive.
Their relative tendency to lose electron in chemical reactions depends on upon:
(i) Ionisation enthalpy: Lower the ionisation enthalpy, greater the chemical reactivity.
(ii) Hydration enthalpy: Larger the hydration enthalpy, greater the chemical reactivity.
Now ionisation enthalpies of alkali metals decrease down the group and hence chemical reactivity increases. Also, hydration enthalpies of alkali metals increase down the group and so chemical reactivity increases. Hence, the chemical reactivity increases down the group of alkali metals. Thus, caesium (Cs) is the most reactive among the alkali metals i.e. Cs has maximum tendency to lose the valence electron to form monovalent Cs⁺ ions.

Lower the ionisation enthalpy, greater is the tendency of an element to lose electrons and hence stronger is the reducing character or higher is the reactivity of the element. Since the ionisation enthalpies of alkali metals decrease down the group, therefore, their reducing character or reactivity in the gaseous state increases from Li to Cs i.e. Li < Na < K < Rb < Cs. However, in the aqueous solution, it has been absorbed that the reducing character of alkali metals follows the sequence Na < K < Rb < Cs < Li. In other words, lithium is the strongest reducing agent in aqueous solution. 

Electrode potential is the measure of the tendency of an element to lose electrons in the aqueous solution. Thus, more negative is the electrode potential, higher is the tendency of the element to lose electrons and hence stronger is the reducing agent.
Since the standard electrode potential  of
alkali metals become more and more negative as we move down the group from Na to Cs, therefore, reducing character of these elements increases in the same order i.e. Na to Cs. However, standard electrode potential (reduction) of lithium is the lowest i.e. -3.05 volts. In other words, lithium is the strongest reducing agent in the aqueous solution. 
Reason. Electrode potential depends on :
(i) heat of sublimation
(ii) ionisation enthalpy
(iii) heat of hydration.

The sublimation enthalpies of alkali metals are almost similar. Now since Li⁺ ion is smallest in size, therefore, the large amount of energy released in step III (heat of hydration) compensates for the higher ionisation enthalpies, thereby facilitating the release of electron and hence explains the low value of electrode potential 

This is because they have a strong tendency to lose the valence s-electron to acquire the nearest inert gas configuration. Lithium, however, because of its high ionisation enthalpy forms covalent compounds i.e. alkyl lithium (R - Li), aryl lithium (Ar - Li) etc. 

It is because they have a strong tendency to lose the valence s-electron to acquire the nearest inert gas configuration. Lithium, however, because of it high ionisation enthalpy forms covalent compounds i.e. alkyl lithium (R – Li), aryl lithium (Ar – Li) etc.
Down the group new shell added which cause shielding effect and thus cause lower the melting point.

Lithium salts are commonly hydrated like LiCl.2H₂O whereas other alkali ions are usually anhydrous. The hydration enthalpy of Li⁺ ion is maximum hydrated and therefore, the effective size of Li⁺ in aqueous solution is the largest. The hydration enthalpy decreases with increase in ionic size.

Li⁺ >Na⁺>K⁺> Rb⁺>Cs⁺

Therefore, the ions of other alkali metals are usually anhydrous.

Lithium is the only alkali metal which forms a nitride directly Li⁺ and N₃^- ions are very small, so they form stable Li₃N.

Lithium forms normal oxide  Lithium-ion with small size has a strong positive field around it. On combination with oxide anion, the positive field of lithium ion restricts the spread of negative charge towards another oxygen atom and thus prevents the formation of a higher oxide. 
Sodium reacts with dioxygen to form sodium peroxide  Sodium ion with a larger size than lithium ion has weaker positive field than lithium ion. This positive field is so weak that it cannot prevent the conversion of the oxide anion  into a peroxide ion , However, it is strong enough to prevent further oxidation of peroxide to superoxide. 
Potassium, rubidium and caesium react with dioxygen to form superoxide 
Potassium, rubidium and caesium ions are large sized and thus have a very weak positive field around them. The positive field around these ions is so weak that it cannot prevent the conversion of peroxide  anion to superoxide anion 

LiF is almost insoluble in water because of much higher lattice energy than that of LiCl. Since Li⁺ ion (small size) can polarise bigger Cl^- ion more easily than the smaller F^- ion, therefore, according to Fajan rules, LiCl has more covalent character than LiF and hence LiCl is soluble in the organic solvents like acetone. 

As the size of Li⁺ ion is much smaller than K⁺ ion, therefore Li⁺ ion can polarise bigger I^- ions to a greater extent than K⁺ ions (Fajan rule). As a result, LiI is more covalent than KI and hence is more soluble in ethanol. (organic solvent).

Anomalous behaviour of lithium is due to its:
(i) very small size,
(ii) high electronegativity and ionisation energy enthalpy value and
(iii) the absence of d-orbitals in the valence shell of its atom.
Therefore, lithium differs from other members of the family in the following respects:
(i) Lithium is harder than other alkali metals.
(ii) Lithium combines with oxygen to form lithium oxide while other alkali metals form peroxides and superoxides.
(iii) Lithium when heated with ammonia forms imide, Li₂NH, while other alkali metals form amides, MNH₂ as:

(iv) Lithium hydroxide and lithium carbonate decompose on heating while the hydroxides and carbonates of other alkali metals do not decompose on heating.

(v) Lithium unlike other alkali metals from no ethynide on reaction with ethyne. 
(vi) Lithium nitrate when heated gives lithium oxide, while other alkali metal nitrates decompose to give the corresponding nitrites. 

Alkali metals combine with halogens to form metal halides, which are ionic crystalline solids having general formula M⁺X^–.

Reactivity of alkali metals with particular halogen increases from Li to Cs. On the other hand, the reactivity of halogens with particular alkali metal M decreases from F₂ to I₂.

Ionic character of alkali metal halides: When a cation approaches an anion, the electron cloud of the anion is attracted towards the cation, thus it gets distorted or polarised. The capacity of the cation to polarise the anion is called polarising power, and the tendency of the anion to become polarised, is known as its polarizability. Now greater the polarisation caused, greater is the neutralisation of charge and consequently the ionic character is decreased (or covalent character is increased). The polarising power of a cation and the polarizability of an anion are determined in term of following.

Fajan’s rules:
(i) Cation’s size: Smaller is the cation, greater is its polarising power, e.g. LiCl is less ionic (or more covalent) than KCl, because the size of Li⁺ ion is much smaller than that of K⁺ ion.
(ii) Anion’s size: Larger is the anion, higher is its polarizability, since the hold on the electron-cloud by the nucleus of anion decreases. Hence, the ionic character of lithium halides is in the order : LiF > LiCl > LiBr > Lil.
Alternatively, the covalent character is in the order Lil > LiBr > LiCl > LiF. Since higher the ionic character, higher is the melting point, consequently melting point of LiF > LiCl > LiBr > Lil.

The similarity between Li and Mg is because of their similar atomic radii (Li = 152 pm; Mg = 160 pm) and ionic radii. (Li⁺ = 76 pm,  Mg²⁺ = 72 pm).
These two elements resemble each other in the following properties:
(i) Both Li and Mg decompose water very slowly with the liberation of hydrogen.
(ii) Both Li and Mg form nitrides  - Li directly and Mg on burning in nitrogen.
(iii) LiOH and Mg(OH)₂ are weak bases. 
(iv) Nitrates of both decompose on heating to give oxides. 
                 
(v) Both LiCl and MgCl₂ are soluble in ethanol. 
(vi) Both LiCl and MgCl₂ are deliquescent and crystalline from aqueous solution as hydrates, LiCl_2. 2H₂O and MgCl₂.8H₂O.

Lithium mainly occurs as silicate minerals but the amount present in any mineral is always small and thus extraction of the metal is not so easy. The chief ores of lithium are:
(i) Spodumene: LiAl(SiO₃)₂ containing 4·6% lithium.
(ii) Triphylite (Li, Na)₃PO₄ (Fe, Mn)₃(PO₄)₂ containing up to 4% lithium.
(iii) Petalite, LiAl (Si₂O₅) containing 2·7 3·7% lithium.
(iv) Lepidolite, (Li, K, Na)₂Al₂(SiO₃)₂ containing 1·5% lithium.
Traces of lithium are also present in milk, blood, plants etc.

It involves the following steps:
(i) Preparation of lithium chloride: The silicate mineral is crushed to fine powder and then boiled with dilute H₂SO₄ and filtered to remove insoluble silica (SiO₂). The mother liquor is then treated with calculated amount of Na₂CO₃ to precipitate aluminium and iron as carbonates which are filtered off. The filtrate is then treated with an excess of Na₂CO₃ to precipitate lithium as Li₂CO₃. The precipitate is dissolved in HCl to form LiCl.
(ii) Electrolysis of lithium chloride. A mixture of dry lithium chloride and potassium chloride is fused and electrolyzed in an electrolytic cell (Down’s cell). Potassium chloride is added to increase the electrical conductivity and also to lower the melting point from 883K to 723K. The cell is operated at a temperature of about 675-775K and voltage of 8-9 volts is applied.
As a result of electrolysis, the following reactions take place:

Chlorine gas is liberated at the anode while molten lithium rises to the surface of the fused electrolyte and collects in the cast iron enclosure surrounding the cathode. The metal thus obtained is 99% pure. 

The lithium metals can not be extracted by the usual procedure due to the following reasons:
(i) The metal can not be obtained by the reduction of its oxide (Li₂O or Na₂O) because it is a strong reducing agent by itself and the common reducing agents such as carbon, hydrogen, magnesium and aluminium cannot be used.
(ii) Metal can not be obtained by the electrolysis of the aqueous solution of its salt like LiCl because the metal formed at the cathode will violently react with water to form lithium hydroxide and hydrogen.
(iii) Even molten LiCl can not be used for the electrolysis because the melting point of the salt is so high that it is quite difficult to attain and maintain this temperature.
(iv) Chlorine, a by-product of electrolysis, will corrode the material of the vessel at this higher temperature.

(i) The ionisation enthalpy of lithium is maximum in the group. Therefore, it can not form monovalent cation (Li⁺) so easily as compared to the other alkali metals.
(ii) According to Fajan rule, Li⁺ ion can polarise I^– ion more than the ion because of the bigger size of the anion. Therefore, lithium iodide has more covalent character than lithium fluoride.

Important uses of lithium:
(i) Lithium is used as deoxidiser in the purification of copper and nickel.
(ii) Lithium is used as a getter or scavenger.
(iii) Lithium bromide is used in medicines as a sedative.
(iv) Lithium carbonate is used in making a special variety of glass.
(v) Lithium aluminium hydride (LiAlH₄) is used as a reducing agent in synthetic organic chemistry.

Occurrence: Sodium does not occur in nature in the free state due to its very reactive nature. In the combined state, it occurs to the extent of about 2.6% in the earth's crust. The main ores of sodium are:
1. Sodium chloride (NaCl) found as rock salt and in sea water. 
2. Sodium nitrite (NaNO₃) or Chile saltpetre.
3. Sodium carbonate (Na₂CO₃)
4. Borax (Na₂B₄O₇ . 10H₂O).

Difficulties in the extraction of sodium: The sodium metal cannot be extracted by the usual procedure due to the following reasons:
1. Sodium is a very strong reducing agent and it cannot be obtained by the reduction of sodium oxide with a reducing agent such as carbon.
2. It cannot be prepared from its aqueous solution by metal displacement method as the liberated metal reacts with water.
3. The metal cannot be isolated by the electrolysis of the aqueous solution of its salt such as NaCl because the sodium formed at cathode at once reacts with water producing sodium hydroxide and liberates hydrogen gas.
The above difficulties are solved by carrying out the electrolysis of sodium chloride containing some calcium chloride and potassium fluoride also in the molten state.
Molten sodium chloride cannot be used due to the following reasons:
(i) The melting point of sodium chloride (1085K) is very high and it is quite difficult to keep the sodium chloride in the molten state at such a high temperature during its electrolysis.
(ii) Since the boiling point of sodium is 1160K, the metal tends to change into vapours at a high temperature of electrolysis.
(iii) At the high temperature of electrolysis, both liberated sodium and molten sodium chloride tend to form a fog which is very difficult to separate.
(iv) The product of electrolysis (sodium and chlorine) can react chemically with the apparatus used for electrolysis at high temperature.
The main function of calcium chloride and potassium fluoride is to lower the melting point of sodium chloride from 1085K to 873K. This is because the above-mentioned difficulties are only due to the high melting point of sodium chloride.

In this method, sodium is prepared by the electrolysis of fused anhydrous sodium chloride. Some calcium chloride and potassium fluoride are added to it. It lowers the melting point of sodium chloride from 1085 to 873K. The electrolysis is carried out in an iron box lined with fire bricks known as Down’s cell.
It consists of a graphite anode projected up through the bottom of the cell which is surrounded by a cylindrical iron cathode. The anode and cathode are separated by a wire gauze shell through which molten sodium chloride can easily pass but melted sodium cannot. The cell is also produced with a storage drum for receiving molten metal, an outlet for the removal of chlorine gas at the top.


On electrolysis, chlorine is liberated at the anode which escapes through the iron hood at the top. Sodium is liberated at the cathode and remains in the wire gauze shell. The level of molten sodium rises and it overflows into the receiver. The following reactions take place:

Sodium metal obtained by this method is about 99.5% pure.

Sodium cannot be kept in the air because its surface gets tarnished due to the formation of a layer of its oxide, hydroxide and carbonate on its surface. 
        
Sodium cannot be kept in water because it reacts with water violently and hydrogen evolved catches fire.
 

Therefore, sodium is kept in kerosene oil.

(i) When sodium is treated with hydrogen halide, hydrogen gas is evolved.
   
(ii) Sodium reacts with acetylene to form sodium acetylide.
                    
(iii) When sodium is heated with hydrogen, sodium hydride is produced. 
         
(iv) When sodium is treated with mercury, sodium amalgam is formed as the product.
            

(i) Due to low ionisation energy, when sodium metal or its salt is heated in a flame, the valence electron is excited to higher energy level. When this excited electron returns to the ground state, the absorbed energy is emitted in the visible region of the electromagnetic spectrum and hence the flame appears coloured. 
(ii) The reducing character of sodium is due to its low ionisation energy as a result of which it can easily lose its valence electrons and act as strong reducing agent. For example,
    

It is used:
(i) in sodium vapour lamps.
(ii) as a coolant in the valves of internal combustion engines and in nuclear reactions.
(iii) in drying organic solvents such as benzene to remove even last traces of moisture.
(iv) as a reagent for the detection of nitrogen, sulphur and halogens in the organic compounds.
(v) in sodium amalgam which is used as a reducing agent in organic synthesis.
(vi) in the preparation of sodium compounds such as Na₂O₂, NaCN, NaNH_2.

(vii) in the manufacture of a sodium-lead alloy which is used in the manufacture of tetraethyl lead which is added to petrol as an antiknock compound.

In this process, brine (i.e. a concentrated solution of NaCl), ammonia and carbondioxide are the raw materials. The chemical reactions involved are:

CO₂ needed for the reaction is obtained by heating calcium carbonate and quick lime (CaO) is dissolved in water to form slaked lime Ca(OH)₂.

NH₃ needed for the reaction is obtained by heating NH₄Cl formed in eq. (1) with Ca(OH)₂ formed in eq. (2).

The only by product of the reaction is calcium chloride (CaCl₂).