CLASS 11th CHAPTER -10 The s Block Elements |Chemistry | NCERT SOLUTION| EDUGROWN

NCERT Exercises

Question 1.
What are the common physical and chemical features of alkali metals ?
Solution.
General characteristics of alkali metals :

1.  Electronic configuration : General electronic configuration of the valence shell is ‘ns¹‘ where ‘n’ gives the number of the outermost shells.
2.  Atomic and ionic sizes : They have the largest size in their respective periods. The atomic size regularly increases upon descending the group.
3.  Ionisation enthalpy : It decreases down the group due to decreasing effective nuclear charge. 2^nd ionisation enthalpy is very high as now an electron has to be removed from stable noble gas configuration.
4.  Melting and boiling points: Due to weak metallic bond strength and large atomic sizes, all alkali metals are soft and have low melting and boiling points.
5.  Metallic character : Alkali metals readily lose their valence electron and form M⁺ ions showing +1 oxidation states. Metallic character increases as we descend.
6.  Flame colouration : They impart characteristic colours to flame.
7.  Photoelectric effect : Due to very low value of ionisation enthalpies, alkali metals exhibit photoelectric effect.
8.  Nature of compounds : Due to highly electropositive character, compounds of alkali metals are ionic in nature.
9.  Reactivity towards air : Their lustrous surface gets tarnished when exposed to air. They burn violently in air or oxygen forming oxides. Li forms monoxide (Li₂O), Na forms peroxide (Na₂O₂) while K, Rb, Cs form superoxides (KO₂, RbO₂, CsO₂).
10.  Reactivity towards water : Alkali metals form hydroxides when they come in contact with water and liberate hydrogen gas.
11.  Reactivity towards hydrogen : They form ionic hydride (M⁺H^–) Ionic character of hydrides increases on descending the group. They act as strong reducing agents.
12.  Reactivity towards halogen : They react with halogens to form ionic halides (M⁺X^–). The ionic character of metal halides increases down the group.
13.  Reducing character : An element which acts as a reducing agent, must have low ionisation energy. Alkali metals act as strong reducing agents as their ionisation energy values are low. Since ionisation energy decreases on moving down from Li to Cs, the reducing property increases in the same order. Thus, Li is weakest reducing agent while Cs is the strongest reducing agent amongst alkali metals in free state.
14.  Solution in liquid ammonia : They dissolve in liq. NH₃ forming blue coloured solutions, which are good reducing agents and good conductors of electricity.

The blue colour of the solution is due to ammoniated electron.

Question 2.
Discuss the general characteristics and gradation in properties of alkaline earth metals.
Solution.

1. Electronic configuration : The valence electronic configuration of atoms of the group II A elements is ns², where ‘n’ is the period number.
2.  Atomic and ionic sizes : The size of the atom increases gradually from Be to Ra. Their ions are also large and size of the ion increases from Be²⁺ to Ra²⁺.
3.  Ionisation enthalpy : The 1^st and 2^nd ionisation energies of these metals decrease from Be to Ba as size increases.
4.  Melting and boiling points : Due to the presence of two electrons in the valence shell and stronger bonding in solid state, they have higher melting and boiling points than corresponding alkali metals.
5.  Metallic character : Due to low ionisation energy values, these metals are highly electropositive and readily form M²⁺ ions.
6.  Flame colouration : Except Be and Mg, other members impart characteristic colours when their salts are introduced in the flame.
7.  Hydration energy : The M²⁺ ions of alkaline earth metals are extensively hydrated to form [M(H₂O)x]²⁺ ions and during hydration a huge amount of energy called hydration energy is released.
   M²⁺ + xH₂O ➝ [M(H₂O)x]²⁺ + Energy
8.  Reactivity towards air or oxygen : They react with air or oxygen slowly on heating. Be, Mg and Ca form normal oxides (MO) while Sr and Ba form superoxides (SrO₂, BaO₂). BeO is amphoteric, MgO is weakly basic and others are distinctly basic.
9.  Reactivity towards water : They have lesser reactivity towards water. Be does not react even with boiling water. Mg forms Mg(OH)₂ liberating H₂ gas with boiling water.
10.  Reactivity towards halogens : All the metals of the group combine with various halogens at appropriate temperature forming halides of the formula MX₂.
11.  Tendency to form complexes : Due to their smaller ionic sizes and greater charge densities Be, Mg metals have highest tendency to form complexes.
    BeF₂ + 2F^– ➝ [BeF₄]^2-
12.  Reducing character : Except Be all other metals of this group are reducing agents.

Question 3.
Why are alkali metals not found in nature ?
Solution.
All the alkali metals have one valence electron, ns¹, outside the noble gas core. The loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. They readily lose an electron to give monovalent M⁺ ions. Thus, due to the reason cited above, alkali metals are never found free in nature but are always found in combined state.

Question 4.
Find out the oxidation state of sodium in Na₂O₂.
Solution.

Question 5.

Explain why is sodium less reactive than potassium.
Solution.
This is mainly due to high ionisation enthalpy of sodium as compared to potassium. Therefore, potassium is more electropositive and a stronger reducing agent than sodium. It also reacts with water more violently than sodium.

Question 6.
Compare the alkali metals and alkaline earth metals with respect to

1.  ionisation enthalpy
2.  basicity of oxides and
3.  solubility of hydroxides.

Solution.

1.  Ionisation Enthalpy: The ionisation energies of alkaline earth elements are higher than those of alkali metals due to higher nuclear charge and smaller radii.
2.  Basicity of oxides : Oxides of alkali metals are stronger bases as compared to those of alkaline earth metals present in the same period, e.g., when Na₂O is dissolved in water, NaOH formed is a stronger base than when MgO is dissolved in water to form Mg(OH)₂. This is due to higher ionisation energies of alkaline earth metals.
3.  Solubility of hydroxides : Alkali metal hydroxides are more soluble in water as compared to the hydroxides of alkaline earth metals present in the same period. This is due to higher lattice energy of the hydroxides of alkaline earth elements as compared to those of alkali metals.

Question 7.
In what ways lithium shows similarities to magnesium in its chemical behaviour?
Solution.

Question 8.
Explain why can alkali and alkaline earth metals not be obtained by chemical reduction methods?
Solution.
Alkali and alkaline earth metals cannot be extracted by the reduction of their oxides and other compounds as they are strong reducing agents themselves and no such reducing agents are there which can reduce them to get pure metal.

Question 9.
Why are potassium and caesium, rather than lithium used in photoelectric cells?
Solution.
The ns¹ electron in K and Cs is so loosely held that even the low energy photons can eject this electron from the metal surface. This property is termed as photoelectric effect. K and Cs are used in photoelectric cells because they have very low ionisation energies. But lithium having very high ionisation energy cannot be used for photoelectric effect.

Question 10.
When an alkali metal dissolves in liquid ammonia the solution can acquire different colours. Explain the reasons for this type of colour change.
Solution.
The alkali metals dissolve in liq. NH3 without evolution of hydrogen. The colour of the dilute solution is blue. The metal atom loses an electron and it combines with ammonia molecule.

The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light and thus, imparts blue colour to the solution. The solutions are paramagnetic and on standing, slowly they liberate hydrogen resulting in the formation of an amide.

In concentrated solution, the blue colour changes to bronze and becomes diamagnetic.

Question 11.
Beryllium and magnesium do not give colour to flame whereas other alkaline earth metals do so. Why ?
Solution.
In the case of Ca, Sr, Ba and Ra, the electrons can be excited by the supply of energy to higher energy levels. When the excited electrons return to the original level, the energy is released in the form of light. In Be and Mg, the electrons are tightly held and excitation by mere flame is rather difficult, thus, they do not show flame colouration. Ca imparts brick red colour, Sr imparts crimson colour, Ba imparts apple green colour and Ra imparts crimson colour to the flame.

Question 12.
Discuss the various reactions that occur in the Solvay process.
Solution.
Solvay process : It is an industrial method for obtaining sodium carbonate from sodium chloride (NaCl) and limestone (CaCO₃). The raw materials required in this process are common salt, ammonia and limestone. The equations for the complete process may be written as :

Sodium hydrogen carbonate crystal separates out and is heated to give sodium carbonate.

Question 13.
Potassium carbonate cannot be prepared by Solvay process. Why ?
Solution.
Solvay process cannot be extended to the manufacture of K₂CO₃ because KHCO₃ is too soluble to be precipitated by the addition of ammonium hydrogen carbonate to a saturated solution of potassium chloride.

Question 14.
Why is Li₂CO₃ decomposed at a lower temperature whereas Na₂CO₃ at higher temperature?
Solution.

Question 15.

Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals.
(a) Nitrates
(b) Carbonates
(c) Sulphates.
Solution.

Question 16.
Starting with sodium chloride how would you proceed to prepare
(i) sodium metal
(ii) sodium hydroxide
(iii) sodium peroxide
(iv) sodium carbonate ?
Solution.

Question 17.
What happens when
(i) magnesium is burnt in air
(ii) quick lime is heated with silica
(iii) chlorine reacts with slaked lime
(iv) calcium nitrate is heated ?
Solution.

Question 18.
Describe two important uses of each of the following:
(i) caustic soda
(ii) sodium carbonate
(iii) quick lime.
Solution.
(i) Caustic soda : It is the commercial name of NaOH. It is used :
(a) in refining of petroleum.
(b) in the manufacture of soap, paper, rayon, drugs and dyes.
(ii) Sodium Carbonate : It is used :
(a) in laundries and in softening of water as washing soda.
(b) in the manufacture of glass, sodium silicate, paper, borax, caustic soda, etc.
(iii) Quick lime : It is used :
(a) in the purification of sugar, manufacture of dye stuffs, bleaching powder, CaC2, mortar, cement, glass, etc.
(b) as a cheap alkali, i.e., as acid neutraliser.

Question 19.
Draw the structure of
(i) BeCl₂ (vapour)
(ii) BeCl₂ (solid).

Solution.

Question 20.
The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain.
Solution.
The solubility of an ionic compound depends on two factors :
(i) lattice energy and
(ii) hydration energy.
These two factors oppose each other.
If lattice energy is high, the ions will be tightly packed in the crystal and therefore, solubility will be low. If the hydration energy is high, the ions will have greater tendency to be hydrated and therefore, solubility will be high. Alkali metals i.e., sodium or potassium hydroxides will be highly soluble in water due to larger size of Na and K as compared to that of Mg and Ca, 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 and carbonates of Na and K are easily soluble in water while the corresponding salts of Mg and Ca are sparingly soluble in water.

Question 21.
Describe the importance of the following :
(i) limestone
(ii) cement
(iii) plaster of Paris.
Solution.
(i) Limestone : It is used,
(a) in the manufacture of quick lime, slaked lime, cement, washing soda and glass,
(b) as a flux in the smelting of iron and lead ores.
(ii) Cement : It is an important building material. It is used in concrete and reinforced concrete, in plastering and in the construction of bridges, dams and buildings.
(iii) Plaster of Paris : The largest use of plaster of Paris is in the building industry as well as plasters. It is used for immobilising the affected part of organ where there is a bone fracture or sprain. It is also employed in dentistry, in ornamental work and for making casts of statues.

Question 22.
Why are lithium salts commonly hydrated and those of the other alkali ions are usually anhydrous?
Solution.
Because of smallest size among alkali metals, Li⁺ can polarise water molecules more easily than the other alkali metal ions and hence get attached to lithium salts as water of crystallisation. For example, lithium chloride crystallises as LiCl.2H₂O but sodium chloride as NaCl.

Question 23.
Why is LiF almost insoluble in water whereas LiCI is soluble not only in water but also in acetone ?
Solution.
Insolubility of LiF in water can be elucidated as follows:
The lithium ion has the highest energy of hydration as it is small in size in comparison to the other alkali metal ions and so it should be highly soluble. But the small Li⁺ and F ions interact very strongly resulting in high lattice energy of LiF which is responsible for its insolubility whereas in LiCl due to the difference in atomic size between Li⁺ and Cl^–, they do not interact very strongly, the lattice energy is comparatively small and the magnitude of hydration enthalpy is quite large. Therefore, LiCI dissolves in water. As LiCl has more covalent character than LiF (Fajan’s rule) thus, it is soluble in organic solvents like acetone.

Question 24.
Explain the significance of sodium, potassium, magnesium and calcium in biological fluids.
Solution.
Biological importance of sodium : Sodium ions are found primarily outside the cells in blood plasma and in the interstitial fluids which surround the cells. These ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells.
Biological importance of potassium : Potassium ions are the most abundant cations within cell fluids, where they activate many enzymes, participate in the oxidation of glucose to produce ATP and with sodium, are responsible for the transmission of nerve signals.
Biological importance of magnesium : All enzymes that utilise ATP in phosphate transfer require magnesium as the cofactor. The main pigment for the absorption of light in plants is chlorophyll which contains magnesium.
Biological importance of calcium: About 99% of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function, maintenance of a regular heartbeat, various aspects of metabolism and blood clotting. The calcium concentration in plasma is maintained by two hormones : calcitonin and parathyroid hormone.

Question 25.
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.

Question 26.
Comment on each of the following observations:

Solution.
(a) This is attributed to the hydration of the cation in water. As a result, size of the cation increases and its mobility decreases. Due to the smallest size, Li* ion is hydrated to the maximum and has least mobility while Cs+ ion due to least hydration has maximum mobility.
(b) Lithium is a very strong reducing agent. As a result, it directly combines with nitrogen to form its nitride (Li₃N).
(c) E° of any M²⁺/M electrode depends upon three factors :
(i) enthalpy of vaporisation,
(ii) ionisation enthalpy
(iii) enthalpy of hydration. Since the combined effect of these factors is approximately the same for Ca, Sr and Ba, therefore, their electrode potentials are nealy constant.

Question 27.
State as to why
(a) a solution of Na2C03 is alkaline?
(b) alkali metals are prepared by electrolysis of their fused chlorides?
(c) sodium is found to be more useful than potassium?
Solution.
(a) Sodium carbonate being a salt of strong base (NaOH) and weak acid (H₂CO₃) forms an alkaline solution upon hydrolysis.

(b) Alkali metals are prepared by the electrolysis of their fused chlorides as if aqueous solution of their salts are used for extraction by electrolytic method then hydrogen is discharged at cathode instead of an alkali metal as the discharge potentials of alkali metals are high.
(c) Sodium is relatively more abundant than potassium. At the same time, it is also less reactive and its reactions with other substances can be better controlled.

Question 28.
Write balanced equations for reactions between
(a) Na₂O₂ and water
(b) KO₂ and water
(c) Na₂O and CO₂.
Solution.

Question 29.
How would you explain the following observations?
(i) BeO is almost insoluble but BeSO₄ is soluble in water,
(ii) BaO is soluble but BaSO₄ is insoluble in water,
(iii) LiI is more soluble than KI in ethanol.
Solution.

Question 30.
Which of the alkali metal is having least melting point ?
(a) Na
(b) K
(c) Rb
(d) Cs
Solution.
(d) : Cs ➝ Least melting point.
As the size of metal increases, the metallic bonding decreases and melting point decreases.

Question 31.
Which one of the following alkali metals gives hydrated salts ?
(a) Li
(b) Na
(c) K
(d) Cs
Solution.
(a) : Li ➝ Highest number of hydrated salts Due to small size of Li, it has high charge density that attracts water molecules strongly.

Question 32.
Which one of the alkaline earth metal carbonates is thermally the most stable ?
(a) MgCO₃
(b) CaCO₃
(c) SrCO₃
(d) BaCO₃
Solution.
(d) : BaC0₃ ➝ Most thermally stable carbonate.
 being a bigger anion stabilises a bigger cation more efficiently. Thus, Ba²⁺ being bigger in size is the most stable.