The S - Block Elements

Alkali metals are first group metals in periodic table.

Common physical properties of alkali metals are:

1) Alkali metals have one electron in there valence shell so are ready to give it easily to complete their octet.

2) They exhibit +1 oxidation state in their compounds.

3) They impart colours to flames. This is because the heat from flame excites the electron present in outer orbital and when it comes back it gives out light in visible spectrum.

4) Alkali metals are soft and can be easily cut. Sodium (Na) can be cut easily by knife.

5) They have low melting points as metallic bond present in metals is quite weak.

6) They show photoelectric effect. When metals like Cs and K are exposed to sunlight, they lose electrons.

Common Chemical properties of alkali metals are:

1) Due to only one electron in their outer shell, they can donate it easily. So they are quite reactive. Sodium(Na) is stored in kerosene instead of water.

2) They react to water to form hydroxides which are generally basic in nature.

2M + H₂O → 2M⁺ + 2OH^- + H₂

3) They react with dihydrogen to form metal hydrides.

2m + H₂→ 2M⁺H^-

4) They are strong reducing agents and as you go down the group their reducing power increases.

General properties of alkaline earth metals are as follow:

Physical Properties:

1) As we go down the group, atomic size increases due to addition of new shells.

2) These metals lose two electrons to acquire noble gas configuration. Therefore, their oxidation state is +2.

3) alkaline earth metals have large size, their ionization enthalpies are found to be fairly low. However, their first ionization enthalpies are higher than the corresponding group 1 metals.

4) As we go down the group, ionisation energy decreases, as it is easy to take out electrons now from shells.

5) Ca, Sr, and Ba impart characteristic colours to flames.

Ca – Brick red

Sr – Crimson red

Ba – Apple green

Chemical properties:

1) Alkaline earth metals react with halogens at high temperatures to form halides.

M +X₂→ MX₂

2) All the alkaline earth metals, except Be, react with hydrogen to form hydrides.

3) They react readily with acids to form salts and liberate hydrogen gas.

M + HCl → MCl₂ + H₂

Alkali metals include lithium, sodium, potassium, rubidium, Caesium, and francium. Alkali metals are not found in nature alone because they are too reactive and thus combine with all elements except noble gases.

In Na₂O₂, oxidation state of Na is +1. O has oxidation state -1 because it exists as peroxide.

Proof: Let us assume oxidation state of Na is x.

Now 2x + 2(-1) = 0

By solving we will get x = +1.

As potassium (K) has more number of orbits, valance electron is held loosely by nucleus. Thus, it is easy to take out that electron. In general rule as we go down the group, Ionisation Enthalpy decreases and thus it becomes easy to take out electrons and hence element becomes more reactive. Ionisation Enthalpy of potassium is 419 KJ/mol^–1 whereas of sodium is 496 KJ/mol^–1.

Ionisation Enthalpy: Due to higher nuclear charge in alkaline earth metals, Electron in alkali earth metals are bounded tightly and thus more energy is required to take them out. Thus, Alkaline Earth metals have more Ionisation Enthalpy.

Basicity of Oxides: Alkali metals oxides are generally more basic than alkaline earth metals.

Solubility of hydroxides: Due to lower ionisation enthalpy, alkali metals are more electropositive and thus are more soluble in water.

Lithium and Magnesium have similarities in chemical behaviour and this relation is known as

‘DIAGONAL RELATIONSHIP’. Some similarities are:

1) Both form monoxides with O₂.

2LiOH → Li₂O + H₂O

Mg(OH)₂ → MgO + H₂O

2) Both form nitrides with Nitrogen.

6Li + N₂→ 2Li₃N

3Mg + N₂→ Mg₃N₂

3) Both elements have tendency to form covalent compounds.

4) Both form complex compounds.

5) Neither Li nor Mg form peroxides or superoxide.

6) Li and Mg do not form solid bicarbonates.

7) Both LiCl and MgCl₂ are soluble in ethanol owing to their covalent nature.

In the process of chemical reduction, oxides of metals are reduced using a stronger reducing agent. Alkali Earth metals are themselves reducing agents. They can’t be obtained by chemical reduction process because no better reducing agents are known than alkali earth metals.

All the three, lithium, potassium, and caesium, are alkali metals. Still, K and Cs are used in the photoelectric cell and not Li. Potassium and Caesium can emit their electrons much easily than Lithium as they have very low ionisation enthalpy, they are used in photoelectric cells rather than Lithium. Ionisation Enthalpy of both elements is much lower than Lithium, so they can emit electrons easily in sunlight too.

When an alkali metal is dissolved in liquid ammonia, it results in the formation of a deep blue coloured solution.These solution of alkali metal in ammonia contains ammoniated cations and ammoniated electrons as shown below:
M + ( x + y ) NH₃→ M+(NH₃)x + e-(NH₃)y
The blue colour of the solution is imparted by ammoniated electrons which absorb energy corresponding to red region of the visible light for the their excitation to higher energy levels and transmit blue colour.

On Heating, the electrons get excited to higher energy level due to taking in of energy. When these electrons return to ground state they emit the extra energy in form if radiations, which are, are in the visibility spectrum hence we can see a flame colour. However, no characteristic colour is shown by Magnesium or Beryllium as their atoms have a smaller size and therefore greater amount of energy is needed for excitation of electron to higher energy level, which is not provided by the flame.

The Solvay process is the major industrial process for production of sodium carbonate. When carbon dioxide gas is bubbled through a brine solution saturated with ammonia, sodium hydrogen carbonate is formed. This sodium hydrogen carbonate is then converted to sodium carbonate.

Various reactions of Solvay process are:

Step 1: CO₂ needed for reaction can be obtained by heating calcium carbonate and quick lime (CaO) is dissolved in water.

CaCO₃ (s) → CO₂ + CaO (II)

Step 2: Carbon dioxide is reacted with this ammonium hydroxide and Sodium Chloride to obtain Sodium Bicarbonate:

NH₃ + H₂O → NH₄OH

NaCl + NH₄OH + CO₂ → NaHCO₃ + NH₄Cl

Step 3: The solution containing crystals of NaHCO₃ is filtered to obtain NaHCO₃.

Step 4: NaHCO₃ is heated strongly to convert it into NaHCO₃.

2 NaHCO₃ (s) → Na₂CO₃ + CO₂ + H₂O (used again)

Na₂CO₃ + H₂O → Na₂CO₃.10H₂O

Reaction in one line:

NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ + NH₄Cl

Potassium carbonate is not precipitated when CO₂ is passed through conc. Solution of KCl saturated with ammonia. Thus, potassium carbonate can't be prepared by Solvay’s process.

As we move down the alkali metal group, the electropositive character increases. This causes an increase in the stability of alkali carbonates. However, lithium carbonate is not so stable to heat. This is because lithium carbonate is covalent. Lithium ion, being very small in size, polarizes a large carbonate ion, leading to the formation of more stable lithium oxide.

On Heating:

Li₂CO₃ → Li₂O + CO₂

Carbonates: -i. Solubility: - Carbonates of alkali metals are soluble in water except Li₂CO₃.ii. Thermal stability: - The carbonates of alkali metals except lithium carbonate are stable towards heat. The carbonates of alkaline earth metals and Lithium carbonate decompose on heating to form oxides with the evolution of CO₂.
Li₂CO₃→ Li₂O + CO₂
MgCO₃→ MgO + CO₂
Nitrates:-i. Solubility:-Nitrates of both alkali and alkaline metals are soluble in water due to more hydration energy than lattice energy.ii. Thermal Stability:-Nitrates of alkali metal except Lithium nitrate on heating forms nitrites with the evolution of oxygen.
NaNO₃→ Na₂O + O₂.
LiNO₃ and nitrates of alkaline earth metals on heating form their respective oxides NO₂ and O₂.
2LiNO₃ → Li₂O + 2NO₂ + O₂
2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂Sulphates: -
i. Solubility: - Sulphates of alkali metals are generally soluble in water except Li₂SO₄.ii. Thermal stability: - The Sulphates of both alkali and alkaline earth metals are thermally stable.

Sodium metal is manufactured by electrolysis of fused mixture of NaCl (40%) and CaCl₂ (60%) in Down’s cell at 873K using iron cathode and graphite cathode. Na is liberated at the cathode and Cl₂ is evolved at anode.

At cathode:

At anode: 2Cl^- (melt) → Cl₂ (g) + 2e^-

Note: Down’s cell is an apparatus in which molten NaCl is electrolysed. In the Down’s cell, Carbon(graphite) acts as anode and iron (Fe) acts as cathode. The working of Down’s cell is given below:

Sodium hydroxide is generally prepared commercially by the electrolysis of aqueous NaCl (brine) in Castner-Kellner cellby using mercury cathode and carbon anode.

⇒ Sodium metal which is discharged at the cathode combines with mercury to form sodium amalgam.

⇒ Cl₂ gas is evolved at the anode.

At cathode:

At anode: 2Cl^-→ Cl₂ + 2e^-

The sodium amalgam thus obtained is treated with water to form sodium hydroxide and hydrogen gas.

Note: Caster- Kellner cell is used in the manufacture of sodium hydroxide (NaOH). The working of Caster Kellner process is given below:

Sodium peroxide is obtained by heating sodium in excess of air. The initially formed sodium oxide reacts with more O₂ to form Na₂O₂.

2Na + O₂ → Na₂O

Sodium carbonate is obtained by Solvay Ammonia process. The Solvay ammonia process includes:

⇒ CO₂ is passed through brine (i.e., a concentrated solution of NaCl) saturated with ammonia.

⇒ Sodium bicarbonate (NaHCO₃) is formed which is sparingly soluble and gets precipitated.

⇒ NaHCO₃ on further heating gives sodium carbonate (Na₂CO₃).

Overall reaction:

Note: Solvay ammonia process is used in the manufacture of sodium bicarbonate. The working of Solvay ammonia process is given below:

(i) When magnesium is burnt in air, it forms magnesium oxide (MgO). The reaction taking place is given below:

2Mg (s) + O₂ 2MgO

(ii) When quick lime (CaO) is heated with silica (SiO₂), it forms calcium silicate (CaSiO₃). The reaction taking place is given below:

(iii) When chlorine reacts with slaked lime (Ca(OH)₂), it forms calcium hypochlorite (Ca(OCl)₂) and calcium chloride (CaCl₂). Both products are constituents of bleaching powder. The reaction taking place is given below:

(iv) When calcium nitrate (Ca(NO₃)₂) is heated, it forms quick lime (CaO), nitrogen dioxide (NO₂) and oxygen. The reaction taking place is given below:

(i) Two important uses of caustic soda (sodium hydroxide) are given below:

⇒ It is used in the manufacture of soap, paper, artificial silk, etc. and in petroleum refining and purification of bauxite.

⇒ It is used in the textile industries for mercerising cotton fabrics.

(ii) Two important uses of sodium carbonate are given below:

⇒ It is used in water softening, laundering and cleaning.

⇒ It is used in the manufacture of glass, soap, borax, etc. and in paper, paints and textile industries.

(ii) Two important uses of quick lime are given below:

⇒ It is used in the manufacture of sodium carbonate from caustic soda (sodium hydroxide).

⇒ It is employed in the purification of sugar and in the manufacture of dyestuffs.

(i) In the vapour state, BeCl₂ exists as a chlorobridged dimer.

This chlorobridged dimer dissociates into the linear monomer at high temperatures (1200K)

(ii) In the vapour state, BeCl₂ has polymeric structure with chlorobridges. The structure is given below:

The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water.

Explanation:

⇒ Due to larger size of Na and K as compared to that of magnesium and calcium, the lattice energies of hydroxides and carbonates of sodium and potassium are much lower than those of hydroxides and carbonates of magnesium and calcium.

⇒ As a result, the hydroxides of Na and K are easily soluble in water while that of corresponding salts of magnesium and calcium are sparingly soluble in water.

Note: Lattice Energy is defined as the amount of energy released when cations and anions are brought from infinity to their respective equilibrium sites in the crystal lattice to form a compound. Lower the lattice energy of a compound, higher will be the solubility.

(i) Importance of Limestone (CaCO₃)

⇒ Specially precipitated, CaCO₃ is extensively used in the manufacture of high quality paper.

⇒ It is also used in as an antacid, mild abrasive in toothpaste.

⇒ It is a constituent of chewing gum as a filler in cosmetics.

(ii) Importance of cement

⇒ Cement is an important building material.

⇒ It is used in concrete and in plastering.

⇒ It is used in the construction of bridges, dams and buildings.

(iii) Importance of Plaster of Paris

⇒ Plaster of Paris is extensively used in the building industry as well as plasters.

⇒ It is used in dentistry, in ornamental work and for taking the casts of statues and busts.

⇒ It is also used for immobilising (prevent from moving) the affected part of organ where there is bone fraction or sprain.

Lithium salts are commonly hydrated and those of the other alkali ions usually anhydrous.

Explanation:

The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

⇒ Because of smallest size among alkali metals, Li⁺ can polarise water molecules more easily than the other alkali metal ions.

⇒ As a result, water molecules get attached to lithium salts as water of crystallisation.

⇒ Hence, Li⁺ has maximum degree of hydration enthlpy and for this reason lithium salts are mostly hydrated.

⇒ For example, lithium chloride crystallises as LiCl 2H₂O

Note: Hydration enthalpy is defined as the energy released when new bonds formation takes place between the ions and the water molecules.

LiF is almost insoluble in water whereas LiCl soluble not only in water but also in acetone.

Explanation:

⇒ The solubility of a compound in water depends on the balance between lattice energy and hydration energy.

⇒ Difference in hydration energy and lattice energy of LiCl:

-876 kJmol^-1 – (-845 kJmol^-1) = -31 kJ mol^-1

⇒ Difference in hydration energy and lattice energy of LiF:

-1019 kJmol^-1 – (-1005 kJmol^-1) = -14 kJ mol^-1

⇒ As a result, difference in lattice energy and hydration energy of LiCl is higher than that of LiF. Hence, LiF is sparingly soluble in water while LiCl is soluble.

⇒ One more reason is that higher the lattice energy, lower will be the solubility. LiF is almost insoluble in water because of much higher lattice energy (-1005 kJ mol^-1) than that of LiCl (-845 kJ mol^-1).

⇒ Furthermore, Li⁺ ion can polarise bigger Cl^-ion more easily than the smaller F^-ion.

⇒ As a result, according to Fajan’s rules, LiCl has more covalent character than LiF and hence is highly soluble in organic solvents like acetone.

Significance of sodium:

⇒ Sodium ions are found primarily on the outside of cells, being located in blood plasma and in the interstitial fluid which surrounds the cells.

⇒ The ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells.

Significance of potassium:

⇒ Potassium ions within the cells activate the enzymes and participate in the oxidation of glucose to produce ATP.

⇒ It combines with sodium and are responsible for the transmission of nerve signals.

Significance of magnesium:

⇒ Magnesium is required as the cofactor in all the enzymes that utilise ATP in phosphate transfer.

⇒ Magnesium is the main pigment for the adsorption of light in plants.

Significance of calcium:

⇒ About 99 % of body, calcium is present in bones and teeth.

⇒ It also plays important roles in neuromuscular function, interneuronal transmission, cell membrane integrity and blood coagulation.

(i) When sodium metal is dropped in water, H₂ gas is evolved which catches fire due to the exothermicity of the reaction.

2Na (s) + 2H₂O (l) → 2NaOH (aq) + H₂ (g)

(ii) When sodium metal is heated in free supply of air, sodium peroxide (Na₂O₂) along with a small amount of sodium oxide (Na₂O) is formed.

(iii) When sodium peroxide dissolves in water, sodium hydroxide (NaOH) and hydrogen peroxide (H₂O₂) is formed.

As we move down the group, the ionic and atomic size of the metals increase. The given alkali metal ions can be arranged in the increasing order of their ionic sizes as:

Li⁺ < N⁺ < K⁺ < Rb⁺ < Cs⁺

Smaller the size of the ion, more highly it is hydrated. Since Li⁺ is the smallest in size, it becomes hydrated in the solution. On the other hand, Cs⁺ is the largest in size and so it is the least hydrated. Hence, the extent of hydration decreases in the order:

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Greater is the mass of the hydrated ion, lower will be its ionic mobility. Hence,Li⁺ ionhas lower ionic mobility and Cs⁺has higher ionic mobility. Thus, the given alkali metal ions can be arranged in the increasing order of their mobilities as:

Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Lithium is the only alkali metal to form a nitride directly. This is because of the following reasons:

⇒ Li⁺ is very small in size and so its size is the most compatible with the N^3– ion.

⇒ One more reason is that because of the diagonal relationship of Li and Mg. Lithium like magnesium forms a nitride while other alkali metals do not.

6Li (s) + N₂ (g) 2Li₃N (s)

E^° of any M²⁺/ M electrode depends upon the following three factors:

⇒ Enthalpy of vaporization

⇒ Ionization enthalpy

⇒ Enthalpy of hydration

Since, the combining effect of three factors is approximately the same for Ca, Sr ad Ba, therefore, their electrode potentials are nearly constant.

A solution of Na₂CO₃ is alkaline because of the following reasons:

⇒ Na₂CO₃ is a salt of weak acid, carbonic acid (H₂CO₃) and a strong base, sodium hydroxide (NaOH)

⇒ Therefore, it undergoes hydrolysis to produce strong base and hence its aqueous solution is alkaline in nature.

Alkali metals are prepared by electrolysis of their fused chlorides because of the following reasons:

⇒ The discharge potential of alkali metals is much higher than that of hydrogen.

⇒ Therefore, when the aqueous solution of any alkali metal chloride is subjected to electrolysis, H₂ instead of the alkali metal is produced at the cathode.

⇒ Therefore, to prepare alkali metals, electrolysis of their fused chlorides is carried out.

⇒ One more reason is that we cannot prepare alkali metals by the chemical reduction of their oxides because alkali metals are very strong reducing agents.

⇒ Even, they cannot be prepared by displacement reactions (one element is displaced by another). This is because these elements are highly electropositive (to lose electrons).

⇒ Neither can electrolysis of aqueous solutions be used to extract these elements. This is because the liberated metals react with water.

⇒ Hence, to avoid these difficulties, alkali metals are usually prepared by the electrolysis of their fused chlorides.

Sodium is found to be more useful than potassium because of the following reasons:

⇒ Sodium ions are found primarily in the blood plasma and in the interstitial fluid which surrounds the cells while potassium ions are present within the cells.

⇒ Sodium ions help in the transmission of nerve signals, in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into the cells.

⇒ Thus, sodium is found to be more useful than potassium.

a) The balanced equations for reactions between Na₂O₂ and water is given below:

Na₂O₂ (s) + 2H₂O (l) → 2NaOH (aq) + H₂O₂ (aq)

b) The balanced equations for reactions between KO₂ and water is given below:

2KO₂ (s) + 2H₂O (l) → 2KOH (aq) + H₂O₂ (aq) + O₂ (g)

c) The balanced equations for reactions between Na₂O and CO₂.is given below:

Na₂O + CO₂→ Na₂CO₃

(i) BeO is almost insoluble but BeSO₄ is soluble in water due to the following reasons:

⇒ Because of the smaller size , higher ionization enthalpy and higher electronegativity, BeO is essentially covalent and hence is insoluble in water.

⇒ On the other hand, BeSO₄ is ionic.

⇒ As we know that, the solubility of a ion in water depends on the balance between lattice energy and hydration energy. The hydration energy of BeSO₄ is much higher than its lattice energy and BeSO₄ is soluble in water.

(ii) BaO is soluble but BaSO₄ is insoluble in water because of the following reasons:

⇒ Both BaO and BaSO₄ are ionic compounds. However, the size of O^2- ion is much smaller than that of SO₄^2- ion.

⇒ Since a bigger anion stabilizes a bigger cation more than a smaller anion stabilizes a bigger cation, therefore, the lattice energy of BaO is much smaller than that of BaSO₄.

⇒ As we know that smaller the lattice energy, greater will be the solubility. Hence, BaO is soluble while BaSO₄ is insoluble in water.

(iii) LiI is more soluble than KI in ethanol because of the following reasons:

⇒ Li⁺ is much smaller than K⁺ ion.

⇒ Therefore, according to the Fajan’s rule, Li⁺ ion can polarise bigger I^- ion to a greater extent than K⁺ ion.

⇒ As a result, LiI is more covalent than KI and hence more soluble in organic solvents like ethanol.

⇒ The size of alkali metal increases as we move down the alkali group.

⇒ Thus as the alkali metal size increases, the binding energies

of their atoms in the crystal lattice decreases. Hence, on moving down the alkali group, the strength of metallic bonding decreases. Because of this, melting point decreases.

⇒ Among the given metals, since the size of Cs is the biggest, therefore, its melting point is the lowest. Thus, option (D) is correct.

⇒ As we know that smaller the size of an ion is, more highly it is hydrated.

⇒ Among alkali metal ions, Li⁺ is the smallest. Therefore, it has the high charge density and high polarising power.

⇒ Li⁺ion attracts the water molecules than any other alkali metal cation.

⇒ Thus, option (A) is correct.

⇒ The thermal stability increases with increase in the electropositive character of the metal or the size of the cation present in the carbonate increases.

⇒ The increasing order of the size of cations of the given carbonates:

Mg < Ca < Sr < Ba

Hence, the increasing order of the thermal stability is:

MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃

⇒ Thus, amongalkaline earth metal carbonates, option (D), BaCO₃ is thermally the most stable.