S Block elements, Class 11, Chemistry Detailed Chapter Notes - JEE PDF Download

S-block Elements : 

The elements in which the last electron enters the outermost s-orbital are called s-block elements. The group 1 and 2 of periodic table belong to the s-block.

Group-I Elements (Alkali Metals) :

(1) The Elements : are Li, Na, K, Rb, Cs Fr (Radioactive: t_1/2 of Fr²³³ = 21 minutes)

group-I elements are called alkali metals because they form hydroxides on reaction with water, which are alkaline in nature.

(2) Outer Electronic configuration : ns¹                                                     

(3) Atomic and Ionic radii :
Li < Na < K < Rb < Cs.
Increase down the group, because value of n(principal quantum number) increases.

(4) Density :
Li < K < Na < Rb < Cs.

(5) Ionization Energy :
Li > Na > K > Rb > Cs.
As size increases, I. E. decreases down the group (so Cs have lowest I. P. )

(6) Hardness and melting points/boiling points :
These metals are very soft and can be cut with a knife. Litium is harder than any other alkali metal.
The hardness depends upon cohesive energy
M. P. Li > Na > K > Rb > Cs
B. P. Li > Na > K > Cs > Rb

(7) Electropositive character or metallic character :
Alkali metals are strongly electropositive and metallic. Down the group electropositive nature increase so metallic nature also increases.
i.e., M → M e^-

Metallic Nature : Electropositive character a  
Li < Na < K < Rb < Cs.

(8) Oxidation state :
Show 1 oxidation state because by loosing one electron they get stable noble gas configuration.

(9) Photoelectric effect :
The phenomenon of emission of electrons when electromagnetic rays strikes against them is called photoelectric effect ; Alkali metal have low I. P. so show photoelectric effect.
*Cs and K are used in photoelectric cells.

Chemical Properties :

(1) Reactions with air
The alkali metals tarnish in dry air due to the formation of their oxides on their surface, which in turn react with water to form hydroxides
4M + O₂ → 2M₂O
M₂O + H₂O → 2MO
They react vigorously in oxygen forming following oxides.
4 Li + O₂ → 2 Li₂O (Monoxide)
2 Na + O₂ → Na₂O₂(Peroxide)
M + O₂ → MO₂(Superoxide) where M = K, Rb, Cs

(2) Solution in liquid NH₃
Alkali metals dissolve in liquid ammonia (high conc. 5 M) and give blue solution which is conducting, reducing and paramagnetic in nature.
Reason
On dissolving Metal in NH₃
M(s)  M   e^-
M  + x(NH₃) → [M(NH₃)_x]  Ammmoniated cation
e^- + y(NH₃) → [e(NH₃)_y]^- Ammmoniated electron
The blue colour is due to → Ammoniated electron
The paramagnetic nature is due to → Ammoniated electron
The conducting nature is due to → Ammoniated M   Ammoniated electron
* On standing the colour fades due to formation of amide
M +  e^- + NH₃ → MNH_2(amide) +  H_2(g)
In the absence of impurities like. Fe, Pt, Zn etc, the solutions are stable.
* In concentrated solution, the blue colour changes to bronze colour and diamagnetic due to the formation of metal clusters and ammoniated electrons also associate to form electron pairs
2e^-(NH₃)y → [e^_(NH₃)₄]₂

(3) Reducing Nature
Reducing agent is electron donor.
Alkali metals are strong reducing agents with lithium being the strongest and sodium the least prowerful reducing agent. Na < K < Rb < Cs < Li
Note : Lithium is expected to be least reducing agent due to it's very high IE. However it is strongest. (due to high hydration energy).

(4) Reaction with H₂O
The reaction with water to form the hdyroxides having the formula MOH
2M + 2H₂O → 2MOH H₂
(Highly reactive)

(5) Reaction with H₂
They react with H₂ forming metal hydride with formula MH which are of ionic nature. Stability of hydride decreases down the group.

(6) Reaction with N₂
Only Lithium reacts with N₂ to form ionic lithium nitride Li₃N.
3Li + N₂ → Li₃N

(7) Reaction with halogens X₂
The alkali metals react vagariously with halogens to form ionic halides M X^-
2M + X₂ → 2MX

(8) Sulphides
All metals react with S forming sulphides such as Na₂S and Na₂S_n(n = 2,3,4,5 or 6). The polysulphide ions are made from zig-zag chains of sulphur atoms.

(9) Crown Ethers and Cryptands

 
Dibenzo-18-Grown-6 Cryptand-222
[Na(Cryptand 222)] Na^- [Contains Na^_(sodide ion)]
[(s (Cryptand-222)] [(Cyrptand-222)e^-] [electride]

Group II Elements (Alkaline earth Metals) :
(1) Elements : The Elements are Be, Mg, Ca, Sr, Ba, Ra,

(2) Outermost Electronic configuration : ns² ,

(3) Atomic and ionic sizes :
* The atomic and ionic radii of the alkali earth metal are smaller than corresponding alkali metals
Reason : higher nuclear charge (Zeff)
* On moving down the group size increase, as value of n increases.
Be < Mg < Ca < Sr < Ba

(4) Ionization Enthalpy :
Be > Mg > Ca > Sr > Ba
Down the group IE decreases due to increase in size
Q. IE₁ of AM < IE₁ of AEM
IE₂ of AM > IE₂ of AEM
[Where AM = Alkali metal, AEM = Alkaline earth metal]
Reason : IE₁ to AEN is large due to increased nuclear charge in AEM as compared to AM but IE₂ of AM is large because second electron in AM is to be removed from action which has already acquired noble gas configuration.

(5) Melting and Boiling points :
They have low m.p. and b.p. but are higher than corresponding value of group I.
Reason : They have two valency electrons which may participate in metallic bonding compared with only one electron in AM. Consequently group II elements are harder and have higher cohesive energy and have much higher m.p./b.p. than A. M.
M.P. Be > Ca > Sr > Ba > Mg
B.P. Be > Mg > Ca > Ba > Sr

(6) Electropositive and Metallic character :

Due to low IE they are strong electropositive but not as strong as AM because of comparatively high IE.
The electropositive character increase down the group.
Be < Mg < Ca < Sr < Ba

(7) Oxidation state : Show 2 oxidation state.

Chemical Properties :
(1) Reactivity towards Air or Oxygen
Be and Mg are kinetically inert towards oxygen because of formation of a film of oxide on their surface.
However powdered Be burn brilliantly.
2 Be + O₂(air)  2BeO
3 Be + N₂(air)  Be₃N₂
Only Mg give the follwoing behaviour.
Mg + Air(N₂ + O₂)  MgO + Mg₃N₂
(Similar property with Li due to diagonal relation.)
* BeO, MgO are used as refractory, because these have high m.p.
* Other metals (Ba or Sr form peroxide)
M + O₂  MO₂

(2) Reaction with H₂O
AEM have lesser tendency to react with water as compared to AM. They form hydroxides and liberate H₂ on reaction with H₂O.
M + 2H₂O  M(OH)₂ + H₂
* Be is inert towards water.
* Magnesium react as
Mg + 2H₂O → Mg(OH)₂  + H₂
or Mg + H₂O → MgO + H₂O
MgO forms protective layer, that is why it does not react readily unless layer is removed amalgamating with Hg.
Other metals react quite readily (Ca, Sr, Ba).
Note : Be(OH)₂ is amphoteric but other hydroxides are basic in nature.

(3) Reaction with Acids & Bases
AEM react with acids & liberate H₂
Mg + 2HCl → MgCl₂ + H₂
Be is amphoteric as it also react with NaOH, other metals do not react as they are purely basic.
Be + 2NaOH → Be(OH)₂  [Be(OH)₄]^2_

(4) Tendency to form Complexes

AEM have tendency to form some stable complexes. Among these Be and Mg have maximum tendency due to their small size and high charge density.

BeF₂  + 2F^- →` [BeF₄]^-2
*Chlorophyll contains Mg²  [Photosynthetic pigment in plants]
*Ca²  and Mg²  form complex with EDTA.

(5) Reactivity with NH₃

Like AM, the AEM (only Ca, Sr, Ba) dissolve in by NH₃ to give deep blue -black solutions having ammoniated cations, and ammoniated electrons.

(6) Reaction with Carbon

AEM when heated with carbon from carbides
*Be form Be₂C
*Mg, Ca, Sr, Ba form carbides of the formula MC₂.

Group - I & II :

Oxides :

Sodium Oxide (Na₂O) :

Preparation :
(i) It is obatined by burning sodium at 180ºC in a limited supply of air or oxygen and distilling off the excess of sodium in vacuum.
2Na + O₂  Na₂O
(ii) By heating sodium peroxide, nitrate or nitrite with sodium.
Na₂O₂ + 2Na → 2Na₂O
2NaNO₃ + 10Na → 6Na₂O N₂
2NaNO₂ + 6Na → 4Na₂O N₂
Properties :
(i) It is white amorphous mass.
(ii) It decomposes at 400ºC into sodium peroxide and sodium
2Na₂O  Na₂O₂ + 2Na
(iii) It dissolve violently in water, yielding caustic soda.
Na₂O + H₂O → 2NaOH

Sodium Peroxides (Na₂O₂) :

Preparation : It is formed by heating the metal in excess of air or oxygen at 300º, which is free from moisture and CO₂.
2Na + O₂ → Na₂O₂

Properties :
(i) It is a pale yellow solid, becoming white in air from the formation of a film of NaOH and Na₂CO₃.
(ii) In cold water (~ 0ºC) produces H₂O but at room temperature produces O₂. In ice-cold mineral acids also produces H₂O₂.

Na₂O₂ + 2H₂O  2 NaOH + H₂O₂

2Na₂O₂ + 2H₂O  4 NaOH + O₂

Na₂O₂ + H₂SO₄  Na₂SO₄ + H₂O₂
(iii) It reacts with CO₂, giving sodium carbonate and oxygen and hence its use for purifying air in a confined space e.g. submarine, ill-ventilated room,

2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂

(iv) It is an oxidising agent and oxidises charcoal, CO, NH₃, SO₂

3Na₂O₂ + 2C → 2Na₂CO₃ + 2Na [Deposition of metallic Na]

CO + Na₂O₂ → Na₂CO₃

SO₂ + Na₂O₂ → Na₂SO₄

2NH₃ + 3Na₂O₂ → 6NaOH + N₂

(v) It containes peroxide ion [-O-O-]^-2
Uses :
(i) For preparing H₂O₂, O_2,
(ii) Oxygenating the air in submarines,
(iii) Oxidising agent in the laboratory.

Oxides of Potassium :
                    K₂O,        K₂O₂ ,     K₂O₃,      KO_2,                       KO₃

Colours : White      White      Red         Bright Yellow       Orange Solid

Preparation :
(i) 2KNO₃ + 10 K  6K₂O + N₂
** K₂O  K₂O
(White) (Yellow)
** K₂O + H₂O  2KOH
(ii) 2K + O₂  K₂O₂ [Props : Similar with Na₂O₂]
(iii) Passage of O₂ through a blue solution of K in liquid NH₃ yields oxides K₂O₂(white), K₂O₃ (red) and KO₂ (deep yellow) i.e.

K in liq. NH₃  K₂O₂  K₂O₃  KO₂

                                   white               red               yellow
** KO₂ reacts with H₂O and produces H₂O₂ and O₂ both
2KO₂ + 2H₂O  2KOH + H₂O O₂
KO₃ : KOH O₃ (ozonised oxygen)  KO₃
(Dry powdered)                                        (orange solid)

Magesium Oxide (MgO) :

It is also called magnesia and obtained by heating natural magnesite.

MgCO₃  MgO + CO₂

Properties :
(i) It is white powder

(ii) It's m.p. is 2850ºC. Hence used in manufacture of refractory bricks for furances.

(iii) It is very slightly soluble in water imparting alkaline reaction.

Calcium Oxide (CaO) :
It is commonly called as quick lime or lime and made by decomposing lime stone at high temperature about 1000ºC.
CaCO₃  CaO + CO₂  42000 cal

Properties :
(i) It is white amorphous powder of m.p. 2570ºC.

(ii) It emits intense light (lime light), when heated in oxygen - hydrogen flame.

(iii) It is an basic oxide and combines with some acidic oxide e.g.
CaO + SiO₂  CaSiO₃

CaO + CO₂  CaCO₃
(iv) It combines with water to produce slaked lime.
CaO + H₂O  Ca(OH)₂

Magnesium Peroxide (MgO₂) and Calcium Peroxide (CaO₂) :
These are obtained by passing H₂O₂ in a suspension of Mg(OH)₂ and Ca(OH)₂.
Uses : MgO₂ is used as an antiseptic in tooth paste and as a bleaching agent.

Hydroxides :

Sodium Hydroxides :

Preparation :

(i) Electrolysis of Brine :
NaCl  Na   Cl^-

At anode : 2Cl^-  Cl₂  2e

At cathode : H +  e^-  H
(ii) Caustication of Na₂CO₃ (Gossage's method) :
Na₂CO₃ + Ca(OH)_2 ⇔ 2NaOH + CaCO_{3 (precipitate)}

(suspension)

Since the K_sp(CaCO₃) < K_sp(Ca(OH)₂), the reaction shifts towards right.
Properties :
(i) It is white crystalline, deliquescent, highly corrosive solid.

(ii) It is stable towards heat.

(iii) It's aqueous solution alkaline in nature and soapy in touch.
(iv) NH₄Cl + NaOH  NaCl + NH₃ ­+ H₂O

FeCl₃ + 3NaOH  Fe(OH)₃ +  3NaCl

                                             ^{Brown ppt}

ZnCl₂ + 2NaOH  Zn(OH)₂  +  2NaCl

Zn(OH)₂+  2NaOH  Na₂ZnO₂ + 2H₂O [Same with AlCl₃, SnCl₂, PbCl₂]

soluble.
(v) Acidic and amphoteric oxides gets dissolved easily e.g.

CO₂ + 2NaOH  Na₂CO₃ + H₂O

Al₂O₃ + 2NaOH  2NaAlO₂ + H₂O

(vi) Aluminium and Zn metal gives H₂ from NaOH

2Al + 2NaOH + 2H₂O  3H₂ + 2NaAlO₂
(vii) Several non metals such as P, S, Cl etc. yield a hydride instead of hydrogen e.g.

4P + 3NaOH + 3H₂O  PH₃ + 3NaH₂PO₂ (Disproportionation reaction)

Potassium Hydroxide :

Preparation : Electrolysis of KCl aqueous solution

Properties : Same as NaOH
**(a) It is stronger base comapred to NaOH.
(b) Solubility in water is more compared to NaOH.
(c) In alcohol, NaOH is sparingly soluble but KOH is highly soluble.
(d) As a reagent KOH is less frequently used but in absorption of CO₂, KOH is preferably used compared to NaOH. Because KHCO₃ formed is soluble whereas NaHCO₃ is insoluble and may therefore choke the tubes of apparatus used.

Magnesium Hydroxide : It occurs in nature as the mineral brucite.
Preparation : It can be prepared by adding caustic soda solution to a solution of Mg-sulphate or chloride soltion.
Mg⁺² + 2NaOH  Na₂SO₄ + Mg(OH)₂

Properties :
(i) It can be dried at temperature upto 100ºC only otherwise it breaks into its oxide at higher temperature. Mg(OH)₂  MgO + H₂O

(ii) It is slightly soluble in water imparting alkalinity.

(iii) It dissolves in NH₄Cl solution

Mg(OH)₂ + 2NH₄Cl  MgCl₂ + 2NH₄OH
** Thus, Mg(OH)₂ is not therefore precipitated from a solution of Mg ² ions by NH₄OH in presence of excess of NH₄Cl.

Calcium Hydroxide :

Preparation : By spraying water on quicklime

CaO + H₂O  Ca(OH)₂

Properties :
(i) It is sparingly soluble in water.
(ii) It's solubility in hot water is less than that of cold water. Hence solubility decreases with increase in temperature.
(iii) It readily absorbs CO₂ as used as a test for the gas.
(iv) it is used as a mortar.

[Mortar is a mixture of slaked lime (1 part) and sand (3 parts) made into paste with water. ]

Carbonates :

Sodium Carbonate:

Preparation :

(i) Leblanc Process :
NaCl + H₂SO₄(conc.)  NaHSO₄ + HCl
NaCl + NaHSO₄  Na₂SO₄ + HCl

^{(Salt Cake)}
Na₂SO₄ + 4C  Na₂S + 4CO ­
Na₂S + CaCO₃  Na₂CO₃ + CaS

(ii) Solvay Process :
NH₃ + H₂O + CO₂  NH₄HCO₃
NaCl + NH₄HCO₃  NaHCO₃ + NH₄Cl
2NaHCO₃  Na₂CO₃ + H₂O + CO₂
Properties :
(i) Anhydrous Na₂CO₃ is called as soda ash, which does not decompose on heating but melts at 852ºC
(ii) It forms number of hydrates.
Na₂CO₃.H₂O  Crystal carbonate  Na₂CO₃  moisture in air

Na₂CO₃.7H₂O  ________

Na₂CO₃.10H₂O  Washing soda
(iii) Na₂CO₃ absorbs CO₂ yielding sparingly soluble sodium bicarbonate which can be calcined at 250º to get pure sodium carbonate.

Na₂CO₃ + H₂O + CO₂  2NaHCO₃
(iv) It dissolved in acid with effervescence of CO₂ and causticised by lime to give caustic soda.
Na₂CO₃ + HCl  2NaCl + H₂O + CO₂

Na₂CO₃ + Ca(OH)₂  2NaOH + CaCO₃
Uses : It is widely used in glass making as smelter.

Potassium Carbonate :

By leblance process, it can be prepared but by solvay process it cannot be prepared because KHCO₃ is soluble in water.
Properties : It resembles with Na₂CO₃, m.p. is 900ºC but a mixture of Na₂CO₃ and K₂CO₃ melts at 712ºC.

Uses : It is used in glass manufacturing.

Calcium Carbonate :
It ocurs in nature as marble, limestone, chalk, coral, calcite etc. It is prepared by dissolving marble or limestone in HCl and removing iron and aluminium present, by precipitating with NH₃ and then adding (NH₄)₂CO₃ to the solution.

CaCl₂ + (NH₄)₂CO₃  CaCO₃ + 2NH₄Cl

Properties:
(i) It dissociates above 1000ºC as follows:

CaCO₃  CaO + CO₂
(ii) It dissolves in water containing CO₂ forming Ca(HCO₃)₂ but is precipitated from the solution by boiling.

CaCO₃ + H₂O + CO  Ca(HCO₃)₂

Magnesium carbonate :
It occurs in nature as magnesite, isomorphous with calcite. It is obtained as a white precipitated by adding sodium bicarbonate to a solution of a magnesium salt; but only basic carbonate, called magnesia alba, having the approximate composition MgCO₃, Mg(OH)₂, 3H₂O is precipitated.

Properties : Same with CaCO₃

Bicarbonates :

Sodium bicarbonates :

Preparation : By absorption of CO₂ in Na₂CO₃ solution.

Na₂CO₃ + H₂O + CO₂  2NaHCO₃

^{(> 100ºC sparingly soluble)}
Uses : It is used in medicine and as baking powder.

Potassium bicarbonates :

Preparation : Same NaHCO₃

Properties : Same with NaHCO₃

But It is more alkaline and more soluble in water compared NaHCO₃.

Magenisum bicarbonate :

MgCO₃ + CO₂ +  H₂O  Mg(HCO₃)₂

Calcium bicarbonate :

CaCO₃ + CO₂ + H₂O ⇔ Ca(HCO₃)₂

Chlorides :

Sodium Chloride : Prepared from brine containing 25% NaCl.

Properties:

(i) It is nonhygroscopic but the presence of MgCl₂ in common salt renders it hygroscopic.

(ii) It is used to prepare freezing mixture in laboratory [Ice-common salt mixture is called freezing mixture and temperature goes down to - 23ºC.

(iii) For melting ice and snow on road.

Potassium Chloride : It is also occurs in nature as sylvyne (KCl) or carnalite (2KCl.MgCl_2.6 H₂O)

Uses : It is used as fertiliser.

Magnesium Chloride :

Preparation : By dissolving MgCO₃ in dil. HCl

MgCO₃ + 2HCl  MgCl₂ + H₂O + CO₂

Properties :

(i) It crystallises as hexahydrate. MgCl₂.6H₂O

(ii) It is deliquescent solid.

(iii) This hydrate undergoes hydrolysis as follows:

MgCl₂.6H₂O  Mg(OH)Cl + HCl + 5H₂O

Mg(OH)Cl MgO + HCl
** Hence, Anh. MgCl₂ cannot be prepared by heating this hydrate.
** Because of this formation of HCl. Sea water cannot be used in marine boilers which corrodes the iron body.

(iv) Anhydrous MgCl₂ can be prepared by heating a double salt like. MgCl₂ . NH₄Cl . 6 H₂O as follows:

MgCl₂.NH₄Cl.6H₂O  MgCl₂.NH₄Cl  MgCl₂  + NH₃ + HCl

Sorel Cement : It is a mixture of MgO and MgCl₂ (paste like) which set to hard mass on standing. This is used in dental filling, flooring etc.

Calcium chloroide :

(i) It is the by-product in solvay process.

(ii) It may also be prepared by dissolving the carbonate in HCl .

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
Properties :

(i) It is deliquescent crystals.

(ii) It gets hydrolysed like MgCl₂ hence anhydrous CaCl₂ cannot be prepared.

CaCl₂ + H₂O  CaO + 2HCl

Hence, anh CaCl₂ is prepared by heating CaCl₂.6H₂O in a current of HCl (dry)

(iii) Anh. CaCl₂ is used in drying gases and organic compounds but not NH₃ or alcohol due to the formation of CaCl₂.8NH₃ and CaCl₂.4C₂H₅OH.

Sodium sulphate :
Preparation :

It is formed in the 1^st step of leblanc process by heating common salt with sulphuric acid.

2NaCl + H₂SO₄ → Na₂SO₄ + 2HCl

Thus the salt cake formed is crystallised out from its aqueous solution as Na₂SO₄.10H₂O. This called as Glauber's salt.
** One interesting feature of the solubility of glauber's salt is: when crystallised at below 32.4ºC, then Na₂SO₄. 10H₂O is obtained but above 32.4ºC, Na₂SO₄ (anh.) comes out.
Properties : It is reduced to Na₂S when fused with carbon.

Na₂SO₄ + 4C → Na₂S + 4CO

Uses : It is used in medicine.

Potassium Sulphate :

It occurs in stassfurt potash beds as schonite K₂SO₄. MgSO₄.6H₂O and Kainite, KCl.MgSO₄.3H₂O from which it is obtained by solution in water and crystallisation. It separates from the solution as anhcrystals whereas Na₂SO₄ comes as decahydrate.

Uses : It is used to prepare alumn.

Magnesium Sulphate :

Preparation :

(i) It is obtained by dissolving kieserite. MgSO₄.H₂O in boiling water and then crystallising the solution as a hepta hydrate. i.e. MgSO₄. 7H₂O. It is called as Epsom salt.

(ii) It is also obtained by dissolving magnesite in hot dil. H₂SO₄.

MgCO₃ + H₂SO₄ → MgSO₄ + H₂O CO₂

(iii) Or by dissolving dolomite (CaCO₃, MgCO₃) in hot dil. H₂SO₄ and removing the insoluble CaSO₄ by filtration.

(iv) It is isomorphous with FeSO₄.7H₂O, ZnSO₄.7H₂O

Calcium Sulphate : It occurs as anhydrite CaSO₄ and as the dihydrate CaSO₄. 2H₂O, gypsum, alabaster or satin-spar.

Properties :

(i) Gypsum (CaSO₄.2H₂O)  2CaSO₄.H₂O (Plaster of paris)

                                                     200^oC

                                                (anhydrous) CaSO₄

                                                  Dead burnt. plaster

(ii) Solubility of CaSO₄ at first increases upto a certain point and then decreases with rise of temperature.

(iii) Plaster of paris is used in mould making due to its porous body.