CBSE Class 11 Chemistry – Chapter 10 The sBlock Elements- Study Materials

Subtopics of Class 11 Chemistry Chapter 10 – The s-Block Elements

  1. Group 1 Elements: Alkali Metals
  2. General Characteristics Of The Compounds Of The Alkali Metals
    1. Oxides And Hydroxides
    2. Halides
    3. Salts Of Oxo-acids
  3. Anomalous Properties Of Lithium
  4. Some Important Compounds Of Sodium
  5. Biological Importance Of Sodium And Potassium
  6. Group 2 Elements: Alkaline Earth Metals
  7. General Characteristics Of Compounds Of The Alkaline Earth Metals
  8. Anomalous Behaviour Of Beryllium
  9. Some Important Compounds Of Calcium
  10. Biological Importance Of Magnesium And Calcium.

The s-Block Elements Class 11 Notes Chemistry Chapter 10

• General Electronic Configuration of s-Block Elements
For alkali metals [noble gas] ns1
For alkaline earth metals [noble gas] ns2
• Group 1 Elements: Alkali metals
Electronic Configuration, ns1, where n represents the valence shell.
These elements are called alkali metals because they readily dissolve in water to form soluble hydroxides, which are strongly alkaline in nature.
• Atomic and Ionic Radii
Atomic and ionic radii of alkali metals increase on moving down the group i.e., they increase in size going from Li to Cs. Alkali metals form monovalent cations by losing one valence electron. Thus cationic radius is less as compared to the parent atom.
• Ionization Enthalpy
The ionization enthalpies of the alkali metals are generally low and decrease down the group from Li to Cs.
Reason: Since alkali metals possess large atomic sizes as a result of which the valence s-electron (ns1) can be easly removed. These values decrease down the group because of decrease in the magnitude of the force of attraction with the nucleus on account of increased atomic radii and screening effect.
• Hydration Enthalpy
Smaller the size of the ion, more is its tendency to get hydrated hence more is the hydration enthalpy.
Hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
Li+ > Na+ > K+ > Rb+ > Cs+
• Physical Properties
(i) All the alkali metals are silvery white, soft and light metals.
(ii) They have generally low density which increases down the group.
(iii) They impart colour to an oxidising flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region.
• Chemical Properties of Alkali Metals
(i) Reaction with air:
When exposed to air surface of the alkali metals get tarnished due the formation of oxides and hydroxides.
Alkali metals combine with oxygen upon heating to form different oxides depending upon their nature.
(ii) Reaction with water:
Alkali metals react with water to form hydroxide and dihydrogen
(iii) Reaction with hydrogen:
The alkali metals combine with hydrogen at about 673 K (lithium at 1073 K) to form hydrides.
2M + H2 ————-> 2M+
The ionic character of hydrides increases from Li to Cs.
(iv) Reaction with halogens:
Alkali metals combine with halogens directly to form metal halides.
2M + X2————–> 2MX
They have high melting and boiling points.
Order of reactivity of M:
(v) Reducing nature:
The alkali metals are strong reducing agents. In aqueous solution it has been observed that the reducing character of alkali metals follows the sequence Na < K < Rb < Cs < Li, Li is the strongest while sodium is least powerful reducing agent. This can be explained in terms of electrode potentials (E°). Since the electrode potential of Li is the lowest. Thus Li is the strongest reducing agent.
(vi) Solubility in liquid ammonia:
The alkali metals dissolve in liquid ammonia to give deep blue solution. The solution is conducting in nature.
M+ (x + y) NH3 ———-> [M (NH3) X]+ + [e (NH3) y]
When light falls on the ammoniated electrons, they absorb energy corresponding to red colour and the light which emits from it has blue colour. In concentrated solution colour changes from blue to bronze. The blue solutions are paramagnetic while the concentrated solutions are diamagnetic.
• Uses of Alkali Metals
Uses of Lithium
(i) Lithium is used as deoxidiser in the purification of copper and nickel.
(ii) Lithium is used to make both primary and secondary batteries.
(iii) Lithium hydride is used as source of hydrogen for meteorological purposes.
(iv) Lithium aluminium hydride (LiAlH4) is a good reducing agent.
(v) Lithium carbonate is used in making glass.
Uses of Sodium
(i) Used as sodium amalgum in laboratory (synthesis of organic compounds).
(ii) Sodium is used in sodium vapour lamp.
(iii) In molten state, it is used in nuclear reactors.
(iv) An alloy of sodium-potassium is used in high temperature thermometres.
Uses of Potassium
(i) Salts of potassium are used in fertilizers.
(ii) Used as reducing agent.
Uses of Cesium
(i) In rocket propellent
(ii) In photographic cells.
• Group 2 Elements: Alkaline Earth Metals
Alkaline Earth Metals: They were named alkaline earth metals since they were alkaline in nature like alkali metals oxides and they were found in the earth’s crust.
Example, Be (Beryllium), Ca, Mg, Sr etc.
• Electronic Configuration
Their general electronic configuration is represented as [noble gas] ns2.
• Atomic and Ionic Radii
Atomic and ionic radii of alkaline earth metals one comparatively smaller than alkali metals. Within the group atomic and ionic radii increases with the increase in atomic number. Reason: Because these elements have only two valence electrons and the magnitude of the force of attraction with the nucleus is quite small.
• Ionization Enthalpies
These metals also have low ionization enthalpies due to fairly large size of atoms. As the atomic sizes increase down the group ionization enthalpies are expected to decrease in the same manner.
Due to their small size in comparison to alkali metals first ionization enthalpies of alkaline earth metals is higher than that of alkali metals.
• Hydration Enthalpies
The hydration enthalpies of alkaline earth metal ions are larger than those of the alkali metals. Thus alkaline earth metals have more tendency to become hydrate. The hydration enthalpies decreases down the group since the cationic size increases.
Be2+ > Mg2+ > Ca2+ > Sr2+ > Ba2+
Metallic character: They have strong metallic bonds as compared to the alkali metals in the same period. This is due to the smaller kernel size of alkaline earth metal and two valence electrons present in the outermost shell.
• Physical Properties
(i) They are harder than alkali metals.
(ii) M.P and B.P are higher than the corresponding alkali metals due to their small size.
(iii) The electropositive character increases down the group.
(iv) Except Be and Mg, all these metals impart characteristic colour to the flame.
(v) The alkaline earth metals possess high thermal and electrical conductivity.
• Chemical Properties
1. Reaction with oxygen. Beryllium and magnesium are kinetically inert to oxygen because of the formation of a thin film of oxide on their surface.
Reactivity towards oxygen increases as going down the group.
2. Reaction with water. Since these metals are less electropositive than alkali metals, they are less reactive towards water.
Magnesium reacts with boiling water or steam. Rest of the members reacts even with cold water.
Mg + 2H20 ——-> Mg(OH)2 + H2
Ca + 2H20 ————> Ca(OH)2 + H2
3. Reaction with halogens. They combine with the halogens at appropriate temperature to form corresponding halides MX2.
M + X2 ——–> MX2 (X = F, Cl, Br, I)
Thermal decomposition of (NH4)2 BeF4is used for the preparation of BeF2.
4. Reaction with hydrogen. These metals except Be combine with hydrogen directly upon heating to form metal hydrides.
• General Characteristics of Compounds of Alkaline Earth Metals
Oxides and Hydroxides
(i) The alkaline earth metals bum in oxygen to form MO (monoxide).
(ii) These oxides are very stable to heat.
(iii) BeO is amphoteric in nature while oxides of other elements are ionic.
(iv) Exept BeO, they are basic in nature and react with water to form sparingly soluble hydroxides.
MO + H2O ———-> M(OH)2
(v) Hydroxides of alkaline earth metals are less stable and less basic than alkali metal hydroxides.
(vi) Beryllium hydroxide is amphoteric in nature.
The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, MX2.
They can also be prepared by the action of halogen acids (HX) on metals, metal oxides, metal hydroxides.
M + 2HX ——-> MX2 + H2
MO + 2HX ——> MX2 + H20
M (OH)2 + 2HX —–> MX2 + 2H20
(i) Except beryllium halides, all other halides of alkaline earth metals are ionic in nature.
(ii) Except BeCl2 and MgCl2 other chloride of alkaline earth metals impart characteristic colours to flame.
(iii) The tendency to form halide hydrates decreases down the group.
For example, (MgCl2– 8 H20, CaCl2– 6 H20, SrCl2– 6 H20, BaCl2– 2 H2O)
(iv) BeCl2 has a chain structure in the solid phase as shown below.
In vapour phase the compound exist as a dimer which decomposes at about 1000K to give monomer in which Be atom is in sp hybridisation state.
(i) The sulphates of alkaline earth metals are white solids and quite stable to heat.
(ii) BeS04 and MgS04 are readily soluble in water. Solubility decreases from BeS04 to BaS04.
Reason. Due to greater hydration enthalpies of Be2+ ions and Mg2+ ions they overcome the lattice enthalpy factor. Their sulphates are soluble in water.
Carbonates of alkaline earth metals are thermally unstable and decompose on heating.
• Some Important Compounds of Calcium
(i) In the manufacture of cement, sodium carbonate, calcium carbide etc.
(ii) Used in the purification of sugar.
(iii) In the manufacture of dye stuffs.
(i) It is used in the manufacturing of building material.
(ii) Used in white-wash as a disinfectant.
(iii) Used to detect C02 gas in the laboratory.
(iii) Calcium Carbonate or Limestone (CaC03)
Preparation: Calcium carbonate occurs in nature in different forms like limestone, marble, chalk etc. It can be prepared by passing C02 through slaked lime in limited amount.
Ca(OH)2 + C02 ———> CaC03 + H20
It can also prepared by the reaction of a solution of sodium carbonate with calcium chloride.
CaCl2 + Na2C03 ————> CaC03 + 2NaCl
(i) In the manufacturing of Quick Lime.
(ii) With MgC03 used as flux in the extraction of metals.
(iii) Used as an antacid.
(iv) In the manufacture of high quality paper.
(iv) Calcium Sulphate (Plaster of Paris) CaS04-1/2H20
Preparation: It is obtained when gypsum CaS04– 2 H20 is heated to 393 K
2(CaS04-2H20) ———-> 2(CaS04) . H20 + 3H20
Above 393 K anhydrous CaS04 is formed, which is called ‘dead burnt plaster’.
(i) It is a white atmosphous powder.
(ii) When it is mixed in adequate quantity of water it forms a plastic hard mass within 15 minutes.
(i) Commonly used in making pottery, ceramics etc.
(ii) Used in the surgical bandages for setting the fractured bone or sprain.
(iii) For making statues, ornamental work, decorative material etc.
(v) Cement
Preparation: Prepared by combining a material rich in CaO with other material such as clay, which contains Si02 along with the oxides of aluminium, iron and magnesium.
Important Ingredients of portland cement:
(Ca2Si04) dicalcium silicate 26%
(Ca2SiO4) Tricalcium silicate 51%
(Ca3Al206) Tricalcium Aluminate 11%
In plastering and in construction purposes.
• s-block elements constitute Group I and II elements.
• General electronic configuration of
Group I = [Noble gas] ns1
Group II = [Noble gas] ns2
• Diagonal Relationship
The first three elements of second period (Li, Be, B) show diagonal similarity with the elements (Mg, Al, Si) of third period. Such similarities are termed as diagonal relationship.
• The alkali metals are silvery-white soft metals. They are highly reactive. Their aqueous solutions are strongly alkaline in nature. Their atomic and ionic sizes increase on moving down the group and ionization enthalpies decrease systematically down the group.
• Alkaline earth metals. They are much similar to alkali metals but due to small size some differences are there. Their oxides and hydroxides are less basic than the alkali metals.
• Sodium hydroxide (NaOH) is prepared by the electrolysis of aq NaCl in Castner- Kellner cell.
Slaked lime Ca(OH)2 is formed by the action of quick lime on water.
• Gypsum is CaS04. 2 H20. On heating upto 390 K CaS04/2H20 (plaster of paris) is formed.

CBSE Class 11 Chemistry Chapter-10 Important Questions

1 Marks Questions

1.Why is Group I elements known as the most electropositive element?

Ans.The loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. They readily lose electron to give monovalent M+ ions.

2.Why is lithium salts mostly hydrated?

Ans. Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated eg. LiCl, 2H2O.

3.Why are melting and boiling points of alkali metals low?

Ans. The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them.

4.What do you mean by diagonal relationship in the periodic table?

Ans. The diagonal relationship is due to the similarity in ionic sizes and /or charge / radius ratio of the elements.

5.Why is lithium kept under kerosene oil?

Ans. Because of their high reactivity towards air and water, they are normally kept in kerosene oil.

6.Name the lightest metal.

Ans. Lithium is the Lightest known metal (density 0.534g (em3)

7.Why alkali metal hydroxides are make the strongest bases?

Ans. The alkali metal hydroxides are the strongest of all bases because the dissolve freely in water with evolution of much heat on account of intense hydration.

8.Why are peroxides and super oxides stable in comparison to other oxides? 

Ans. The stability of peroxides and super oxides is due to the stabilization of large anions by larger cat ions through lattice energy effects.

9.Name the anomalous properties of lithium.

Ans. The anomalous behaviors of lithium is due to the following-

(i) Exceptionally small size of its atom and ion., Li+

(ii) High polarizing power (I, e; charge / radius radio)

10.Why are lithium compounds soluble in organic solvents?

Ans.Due to high polarizing power, there is increased covalent character of lithium compounds which is responsible for their solubility in organic solvents.

11.How is sodium carbonate prepared? 

Ans. Sodium carbonate is generally prepared by Solvay’s process.

12.What is sodium amalgam?

Ans. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam.

13. Why is sodium hydrogen carbonate known as baking soda?

Ans.Sodium hydrogen carbonate is known as baking soda because it decomposes on heating to generate bubbles of CO2 (leaving holes in cakes and bread)

14.Why does table salt get wet in rainy season?

Ans. Table salts contains impurities of CaCl2 and MgCl2 which being deliquescent compounds absorbs moisture from the air in rainy reason.

15.What is the formula of soda ash?

Ans. Na2CO3

16.Why do alkaline earth metals have low ionization enthalpy?

Ans. The alkaline earth metals have low ionization enthalpies due to fairly large size of atoms.

17.State one reason for alkaline earth metals in general having a greater tendency to form complexes than alkali metals. 

Ans. Because of small size and high charge, the alkaline earth metals have a tendency to form complexes.

18.Compounds of alkaline earth metals are more extensively hydrated than those of alkali metals. Give reason.

Ans. The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions because of smaller six.

19.The melting and boiling points of alkaline metals are higher than alkali metals. Give reason. 

Ans. The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes.

20.What is the nature of oxide formed by Be?

Ans. BeO is covalent and amphoteric while oxides of other elements are ionic and basic in nature.

21.Why does beryllium show similarities with Al?

Ans. Because of their similarity in charge / radius ratios

22.Why is beryllium carbonate unusually unstable thermally as compared to the other carbonates of this group?

Ans. This is due to strong polarizing effect of small Be2+ on the large CO32- anion and leading to the formation of more stable BeO.

23.Why sulphates of Mg and Be soluble in water?

Ans. The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water.

24.Why beryllium is not attacked by an acid easily?

Ans. Beryllium is not readily attacked by acids because of the presence of an oxide film on the metal.

25.Mention the main compounds which constitute Portland cement.

Ans. The main compounds present in Portland cement are-

(i) Dicalcium silicate  

(ii) Tricalcium silicate 

(iii) Tricalcium aluminate 

26.What happens when gypsum is heated to 390K?

Ans. Plaster of parts is formed

27.Anhydrous calcium sulphate  can not be used as plaster of Paris. Give reason.

Ans.Because it does not have the ability to set like plaster of Paris.

28.Mention the natural sources of calcium carbonate. 

Ans.Calcium carbonate occurs in nature in several forma like limestone, chalk, marble etc.

29.What is milk of lime?

Ans. A suspension of slaked lime in water is known as milk of lime.

30.What happens when CaCO3 is subjected to heat?

Ans. On heating CaCO3, quick lime is obtained

31.Show with an example that Ca O is a basic oxide?

Ans. Ca O combines with acidic oxides at high temperature

6CaO + P4O10 

2 Marks Questions

1.Why are lithium halides covalent in nature?

Ans. Lithium halides are covalent because of the high polarization capability of lithium ion The Li+ ion is very small in size and has high tendency to distort electron cloud around the negative halide ion.

2.What makes lithium show properties different from rest of the alkali metals?

Ans. Lithium is a small atom and it forms smaller Li+. As a result, it has very high charge to radius ratio. This is primarily responsible for the anomalous behavior of lithium.

3.Why do alkali metals and salts impart color to an oxidizing flame?

Ans. This is because the heat from the flame excites the outer orbital electron to a higher energy level.

4.What type of oxide is made by sodium?

Ans. Sodium mostly form peroxide when reacted with oxygen

5. Why is potassium lighter than sodium?

Ans.Potassium is lighter than sodium probably because of an unusual increase in atomic size of potassium.

6.Name the alkali metals that form super oxides when heated in excess of air.

Ans.Potassium, rubidium and caesium form super oxides when heated in excess of air.

7.Write a reaction to show that bigger cat ions stabilize bigger anions.

Ans.In the reaction

The larger Cation K+ stabilizes the larger anion I

8.Lithium shows similarities with magnesium in its chemical behavior. What is the cause of these similarities?

Ans.Due to (diagonal relationship)

(i) Similarity in atomic size

(ii) Similar charge to size ratio.

9.Why metals like potassium and sodium can not be extracted by reduction of their oxides by carbon?

Ans.Potassium and sodium are strong electropositive metals and have great affinity for oxygen than that of carbon. Hence they Cannot be extracted from their oxides by reduction with carbon.

10.Give the important uses of sodium carbonate.


(i) It is used in water softening laundering and cleaning

(ii) It is used in the manufacture of glass, soap, borax and caustic soda.

11.What is the difference between baking soda and baking powder?

Ans.Baking soda is sodium bicarbonate (NaHCO3). Which baking powder is a mixture of sodium bicarbonate (NaHCO3) and potassium hydrogen tartar ate.

12.Discuss the various reactions that occur in the solvay process


13. Give two uses of sodium carbonate?

Ans.(i) It is used in the manufacture of soap, glass, paper, borax and caustic soda etc.

(ii) It is used in textile industry and also in petroleum refining.

14. Solution of Na2 CO3 is alkaline. Give reason.

Ans.The solution of Na2CO3 is alkaline in nature because when Na2CO3 is treated with water, it gets hydrolyzed to form an alkaline solution:

15.Name the elements present in Group 2

Ans. Beryllium, Magnesium, Calcium, Strontium, Barium and Radium.

16.The atomic radii of alkaline earth metals are smaller than those of the corresponding alkali metals. Explain why?

Ans. The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same period because of the increased nuclear charge in there elements.

17.The second ionization enthalpy of calcium is more than the first. How is that calcium forms CaCl2 and not CaCl give reasons.

Ans. The higher value of second ionization enthalpy is more than compensated by the higher enthalpy of hydration of Ca2+. Therefore formation of CaClbecomes more favorable than CaCl energetically.

18.Name the metal amongst alkaline earth metals whose salt do not impart colour to a non-luminous flame.

Ans. Beryllium does not impart colour to a non-luminous flame.

19. Which member of the alkaline earth metals family has:

(i) least reactivity

(ii) lowest density

(iii) highest boiling point

(iv) maximum reduction potential

Ans. (i) Be

(ii) Ca

(iii) Be

(iv) Be

20.The alkaline earth metals are called s – block elements. Give reasons.

Ans. Alkaline earth metals are called s – block elements because the last electron in their electronic configuration occupies the s – orbital of their valence shells.

21.Why is Calcium preferred over sodium to remove last traces of moisture from alcohol?

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

22.Name the metal amongst alkaline earth metals whose salt do not impart colour to a non – luminous flame.

Ans. Beryllium does not impart colour to a non – luminous flame.

23. Give the reaction of magnesium with air?

Ans. Magnesium burns with dazzling brilliance in air to give Mg O and Mg3N2

24.Beryllium is reducing in nature. Why?

Ans. Reducing nature is due to large hydration energy associated with the small size of Be2+ ion and relatively large value of the atomization enthalpy of the metal.

25.Give two uses of

(i) caustic soda

(ii) quick lime

Ans.(i) Caustic soda –

(a) It is used in the manufacture of soap, paper, artificial silk and a number of chemicals.

(b) It is used in petroleum refining and purification of bauxite

(ii) Quick time –

(a) It is used in the manu facture of dye stuffs.

(b) It is used in the manu facture of sodium carbonate from caustic soda.

26.What is quick lime? What happens when we add water to it?

Ans.Ca O is quick lime. When we add water to it slaked limes Ca (OH)2 is formed.

27.What is the formulae of caustic potash?s

Ans. KOH.

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