Hydrogen Chapter 9 Class XI notes

HYDROGEN

Chapter - 9

It is placed in first group and first period (first element of the periodic table). It is the lightest element known. It exists as a diatomic molecule i.e. H₂. That is why it is called dihydrogen.

It was discovered by Henry Cavendish in 1766. He prepared it by the action of dilute H₂SO₄ on iron. Its name hydrogen was proposed by Lavoisier because it produces water on burning with oxygen gas. [Hydro = water, gen = producer]

Unique position of Hydrogen:

a) It resembles alkali metals and can be placed in group 1. It can lose one electron [1s¹] and has electropositive character with +1 oxidation state. It combines with electronegative elements or non-metals like H₂O, HCl, H₂S.

b) It resembles halogens and can be placed in group 17. It can gain one electron [1s¹] and has electronegative character with -1 oxidation state. It combines with metals as NaH, CaH₂.

Occurrence: It is the ninth element in order of abundance (0.9% earth’s crust by weight). Isotopes of Hydrogen are ¹H₁ (protium- most abundant 99.985%), ²H₁ (deuterium 0.015%), ³H₁ (tritium 10^-15%). Deutrium was discovered by Urey in 1934.

Preparation of dihydrogen: 

a) Lab preparation – It is prepared by reaction of granulated zinc with dilute HCl.    

        Zn + 2 HCl (dilute) → ZnCl₂ + H₂

         Zn + 2 NaOH (aq) → Na₂ZnO₂ + H₂

b) Commercial preparation – 

(i) By electrolysis of water:   

2 H₂O (l)   →  2H₂(g) + O₂(g)

Electrolyte = water + acid or alkali (to make it a good conductor), cathode = Fe sheet, anode = Nickel plated iron sheet or platinum electrodes.

(ii) By reaction of steam on hydrocarbons at high temperature:

C_nH_2n+2 + nH₂O →    n CO + (2n+1) H₂

E.g. CH₄ (g) + H₂O (g)   →   CO + 3 H₂   [Mixture of CO and H₂ is known as ‘water gas’. It is used for synthesis of methanol and many hydrocarbons, so it is known as synthesis gas or syn gas.]

(iii) Syn gas or water gas is also prepared from sewage, saw-dust, scrap wood, etc. Process of producing syn-gas from coal is known as coal gasification. 

C(s) + H₂O(g) → CO + H₂(g)

(iv) Water-gas shift reaction or Bosch process –

CO + H₂ + H₂O   →    CO₂ + 2H₂ (g)

OR    

C + 2H₂O(g)  →   CO₂ + 2H₂(g)

Properties of H₂ : 

a) Physical properties – Colourless, odourless, tasteless and combustible gas with high calorific value, lighter than air and insoluble in water.

b) Chemical properties –

(i) It is neutral to litmus paper.

(ii) It has high bond dissociation energy (H-H bond).

(iii) N₂ + 3H₂→ 2NH₃ (Haber’s process).

(iv) Acts as good reducing agent: H₂ + Pd²⁺→ 2H⁺ + Pd.

(v) 2C + H₂    →   HCΞCH (g) (ethyne or acetylene)

Uses: Hydrogen is used to prepare ammonia, used in hydrogenation of vegetable oils, used as rocket fuel in liquid form, used in air-ships and balloons for meteorological purposes.

Hydrides – Compound of an element with hydrogen.

Classification of hydrides:

a) Ionic hydrides or saline or salt like hydrides – These are formed by these metals whose electronegativity values are less than hydrogen (or those which are more electropositive than hydrogen) i.e. s-block elements [ group 1 and 2 metals except Be and Mg].

Properties:

(i)            Ionic hydrides are crystalline, non-volatile and conduct electricity in aqueous solution.

(ii)          They have high melting and boiling point.

(iii)         They act as strong bases  LiH + H₂O → LiOH + H₂.

(iv)         They react violently with water to form H₂. NaH + H₂O → NaOH + H₂.                                         

(v)          4LiH + AlCl₃ → LiAlH₄ + 3LiCl  [LiAlH₄ acts as a strong reducing agent]

b) Covalent hydrides – These are formed by p-block elements and Be and Mg as their electronegativity difference with hydrogen is very small. These hydrides usually consist of covalent molecules which are held together by weak Vander Waal’s force of attraction. So, they are known as covalent or molecular hydrides.

Classification of molecular hydrides-

(i)     Electron deficient hydrides – B₂H₆, BeH₂ -These are lewis acids (electron acceptors).

(ii)    Electron precise compounds – CH₄ -They have complete octets.

(iii) Electron rich hydrides – NH₃, H₂O, HF -They have lone pairs (lewis bases i.e. electron donors).

c) Metallic hydrides – They are formed by d-block and f-block elements except metals of group 6,7,8 and 9 (but CrH chromium hydride also exists). These hydrides have properties similar to those of parent metals. So, they are metallic hydrides (except Ni, Pd).

Hydrogen atoms being small in size, occupy some space in the metallic lattice (interstitial sites) but not all. So, they are interstitial hydrides. E.g. in Pd, Pt. This property of high potential for hydrogen storage is used as a source of energy.

They conduct electricity. They are non-stoichiometric (are deficient in hydrogen) like LaH_2.87, NiH_0.6-0.7, PdH_0.6-0.8 i.e. they do not hold law of constant proportion.

Water – (human body contains 65% of water). It is a universal solvent as it can dissociate almost all compounds in it and has high dielectric constant.

Structure of water –

(i) In gaseous state, water exists as discrete molecule. So, it is a bent molecule with bond angle of 104.5⁰ and O-H bond length of 95.7pm (a highly polar molecule with sp³ hybridisation.  Magnetic moment = μ = 1.84 D as Oxygen is an electronegative atom.

(ii) In liquid state, water molecules are held by intermolecular H-bonds and are in association. Each water molecule is generally H-bonded with four other water molecules. (Dipole-dipole interactions)

(iii) In solid state, crystalline form of water is in hexagonal form at atmospheric pressure but at low temperature, it condenses to cubic form. In structure, each oxygen atom is surrounded tetrahedrally by four other oxygen atoms distance of 276pm. It contains void spaces in its cage-like structure.

Ø                Water has maximum density at 277K or 4⁰C. Reason- Ice contains a large number       of vacant spaces in it due to which it has large volume and lesser density in its            cage like structure. When ice melts, some of the bonds are broken and cage-like           structure partially breaks up. Therefore, volume decreases and density increases.

As temperature is raised further above 273K, more H-bonds break resulting in more decrease in volume and increase in density. It goes on till 4⁰C and all void spaces are broken up.

As temperature is raised above 277K, there is increase in volume due to expansion of liquid water and decrease in density. Hence, density of water is maximum at 277K.

Ø               

             Ice floats on water because density of ice is less than water as the ice has cage like  structure having many void spaces due to which it has more volume and less               density.

 Physical properties of water:

a) Colourless, odourless and tasteless.                                                                    

b) Unusual properties of water – It has high boiling point, freezing point, heat of fusion and vapourization than H₂S, H₂Se, H₂Te in its group due to the presence of intermolecular H-bonding between water molecules. It has high specific heat, surface tension, thermal conductivity, dipole moment and dielectric constant.                                                                                                                       

c) It is an excellent solvent for the transportation of ions and molecules required for plants and animals body.

Q Why H₂O, NH₃ and HF have high melting and boiling points?                                                                     

 Ans. They consist of N, O and F electronegative atoms which result in intermolecular H-bonding between molecules and leads to association. Order of boiling point: H₂O > HF > NH₃.

As water is associated with four other water molecules by H-bonding which HF is associated with two other HF molecules. F is more electronegativity than N. So, HF has more stronger H-bonding than NH₃.

Chemical properties of water:

a) H₂O is very stable at room temperature due to highly negative heat of formation but it decomposes at very high temperature. 2H₂O  →   2H₂ (g) + O₂(g)    

b) It has amphoteric nature. Water is weak electrolyte as it shows self-ionization or auto-protolysis. H₂O+H₂O → H₃O⁺ + OH^-                                                                                                                                        

It acts as both acid and base:   (i) H₂O + HCl → H₃O⁺ + Cl^-                

                                                  (ii) H₂O + NH₃ → NH₄⁺ + OH^- 

c) It undergoes redox reactions: (i) 2Na + H₂O → 2NaOH + H₂ (as oxidizing agent)          

(ii) 2F₂+2H₂O→4HF+O₂ (as reducing agent)                                                                                                                 

d) Hydrolytic reactions – It can hydrolyse many oxides, phosphides, nitrides, etc. due to its high dielectric constant.

CaO + H₂O → Ca(OH)₂ , 

SO₂ + H₂O → H₂SO₃ (sulphurous acid),                                          

SiCl₄ + 2H₂O → SiO₂ + 4HCl,  

P₄O₁₀ + 6H₂O → 4 H₃PO₄ (phosphoric acid),                                                        

Mg₃N₂ + 6H₂O → 2NH₃ + 3Mg(OH)₂       

e) Hydrates formation –Water can be associated with ionic compounds as hydrates. This water in coordination with ionic salts is known as water of crystallisation.

Different types of hydrates:                       

 i) Coordinated water: as a constituent of complex [Cr(H₂O)₆]³⁺.3Cl^-

ii) Interstitial water: water is present in the voids of lattice. E.g. BaCl₂.2H₂O

iii) Hydrogen-bonded water: E.g. CuSO₄.5H₂O as [Cu(H₂O)₄]²⁺SO₄^2-.H₂O

Hard and soft water-

Soft water – Water that produces lather with soap readily being free from soluble salts of calcium and magnesium is called soft water. E.g. rain water, distilled water, demineralised water.

Hard water – Water which does not produce lather with soap readily is called hard water due to the presence of bicarbonates, chlorides and sulphates of Ca and Mg ions in it. Hard water forms curdy white precipitate with soaps. Soap is sodium or potassium salt of higher fatty acids like steric acid, palmitic acid or oleic acid.

2 C₁₇H₃₅COONa + CaCl₂→ (C₁₇H₃₅COO)₂Ca (white ppt or scum) + 2NaCl

So, it is unsuitable for laundry and harmful for boilers as it reduces the efficiency of boilers as lot of soap gets wasted in precipitating out Ca²⁺ and Mg²⁺ ions.

Hardness of water –

Types of hardness:

a) Temporary hardness – due to the presence of Ca and Mg ions in the form of bicarbonates. It can be removed by-

(i) By boiling and filtering –

Mg(HCO₃)₂   →  Mg(OH)₂ (insoluble) + 2CO₂ ;

Ca(HCO₃)₂   →   CaCO₃(insoluble) + CO₂ + H₂O

(ii) By Clark’s process – by adding known amount of lime to hard water.

CaO + H₂O → Ca(OH)­₂ ;                

Ca(HCO₃)₂ + Ca(OH)­₂ → 2CaCO₃ (insoluble) + 2H₂O;

Mg(HCO₃)₂ + Ca(OH)­₂ → CaCO₃ (insoluble) + 2H₂O + Mg(OH)₂.

If excess of lime is added then water would again become hard due to the absorption of CO₂.   Ca(OH)­₂ + CO₂ → Ca(HCO₃)₂.

b) Permanent hardness – due to the presence of soluble chlorides and sulphates of calcium and magnesium ions. It is known as non-carbonate hardness. It can be removed by –

(i) Washing-soda process – by adding known amount of Na₂CO₃ and filtering it.                                          CaCl₂ + Na₂CO₃→CaCO₃ + 2NaCl;     

CaSO₄ + Na₂CO₃→ CaCO₃ + Na₂SO₄.

(ii) Ion-exchange resin method – Ca²⁺ and Mg²⁺ ions are exchanged by the ions present in a complex salt called as zeolites or permutit. E.g. Hydrated sodium aluminium silicates Na₂Al₂Si₂O₈.xH₂O (Na₂Z) (naturally occurring).   

Na₂Z + CaCl₂→CaZ + 2NaCl.

Permutit can be regenerated by treating it with aqueous NaCl.           

CaZ + 2NaCl→Na₂Z + CaCl₂.  

Advantage: This method is efficient, cheap and can be used to remove both temporary and permanent hardness.

(iii) Calgon’s method – Calgon( sodiumhexametaphosphate ) is added to hard water.                                        Na₂P₆O₁₈→ 2Na⁺ + Na₄P₆O₁₈^2-;       

Ca²⁺ + Na₄P₆O₁₈^2-→ [Na₂CaP₆O₁₈]^2- + 2Na⁺

(iv) Synthetic resin method – This method is more efficient than ion-exchange. It involves synthesis of cation-exchangers. Synthetic organic exchangers are called as ion-exchange resins. They remove hardness in water and resulting water is known as demineralised water or de-ionized water.            

Cation-exchange resins contain a large hydrocarbon framework attached to acidic groups such as –COOH and –SO₃H group and they can exchange Na⁺ ions with Ca²⁺ and Mg²⁺.                                                           

RSO₃H + Na⁺→RNa + H⁺ (makes water acidic);       

2RNa + Ca²⁺→ R₂Ca + 2Na⁺.    

Resin can be regenerated back by adding aqueous HCl or aqueous NaCl.                                                        

Anion-exchange reins contain a large hydrocarbon framework attached to basic groups such as OH^- as in the form of substituted NH₄OH ( R⁺-NH₃OH^- )

R⁺-NH₃OH^- + Cl^-→R⁺-NH₃Cl^- + OH^- ;      

R⁺-NH₃OH^- + SO₄^2-→ (R⁺-NH₃)₂SO₄^2- + OH^- (makes water basic)

Resin can be regenerated by adding NaOH.

H₂O₂ = Hydrogen peroxide:

Preparation: (i) Lab preparation – BaO₂.8H₂O + H₂SO₄ → BaSO₄ + H₂O₂ + 8H₂O

(ii) Industrial preparation – by electrolysis of 50% sulphuric acid with platinum anode and graphite cathode. Reactions:

2H₂SO₄→ 2H⁺ + 2HSO₄^-;  

at cathode: 2H⁺ + 2e →H₂ ;            

at anode: 2HSO₄→ H₂S₂O₈ + 2e

H₂S₂O₈ + 2H₂O → 2H₂SO₄ + H₂O₂.

This method can be used to prepare D₂O₂.   

K₂S₂O₈ + 2D₂O → 2KDSO₄ + D₂O₂.

Physical properties of H₂O₂–

(i) It is a thick syrupy liquid with pale blue colour.

(ii) It is bitter in taste and more viscous than H₂O due to highly associated with H-bonds.

(iii) It dissolves in water, alcohol and ether as H₂O₂.H₂O . 

(iv) It has dipole moment  of 2.14D [more than water (1.84D)]

Chemical properties of H₂O₂ :

(a) It is an unstable liquid (auto-oxidation and auto-reduction).

2 H₂O₂  →   2H₂O + O₂

(b) Oxidising character:

(i) acidic medium-  2Fe²⁺ + 2H⁺ + H₂O₂ → 2Fe³⁺ + 2H₂O

(ii) in basic medium-   2Fe²⁺ + H₂O₂ → 2Fe³⁺ + 2OH^-

(c) Reducing character:

(i) basic medium- 2MnO₄^- + 3H₂O₂ → 2MnO₂ + 3O₂ + 2H₂O + 2OH^-

(ii) acidic medium-  Mn²⁺ + H₂O₂ → Mn⁴⁺ + 2OH^-

Structure of H₂O₂:  Non-planar molecule. 

2 Oxygen atoms are linked by single covalent bond and each oxygen is linked to hydrogen atom by single bond.

Storage of H₂O₂: As H₂O₂ easily decomposes so it is stored in plastic vessels or wax-lined glass in dark and urea can be added as stabiliser.     

2 H₂O₂   →    2H₂O + O₂

It is kept away from dust because dust can induce explosive decomposition of H₂O₂.

Uses-

1) H₂O₂ is used for bleaching silk, hair, textile, oils, fats, leather (oxidising agent). 

2) It is used as an antichlor (to remove excess Cl₂) in bleaching in textile industry.  

3) It is used as an antiseptic.

Write short note on Fuel Cell & Hydrogen Economy.

s-block elements Chapter 10 Class XI

s-block elements

Chapter - 10

Representative elements = s + p block elements

s-block : Group 1 and group 2

Group 1 = alkali metals as they form water soluble bases like NaOH, KOH

     = Li, Na, K, Rb, Cs, Fr

General configuration = ns¹,   valency = valence electron = 1 (all metals)

 Group 2 = alkaline earth metals or ground metals

     = Be, Mg, Ca, Sr, Ba, Ra

General configuration = ns²,   valency = valence electron = 2 (all metals).

Occurrence – All are reactive metals and do not exist in free state but are present in combined state as halides, oxides, carbonates, silicates, nitrates and phosphates.

Diagonal Relationship – Some elements of certain groups in second period resemble with the certain elements of next higher period (third period). This is known as diagonal relationship.                                    

Reasons: Similar atomic size, electropositive character, ionization energy, polarising power.

For Eg. (i) Li is diagonally related to Mg

(ii) Be is diagonally related to Al

(iii) B is diagonally related to Si.

Anomalous behaviour of elements of second period –

Reasons:

a) small size,

b) absence of d-orbital in valence shell,

c) high ionization energy,

d) high polarisation power of its cation.

Alkali metals - Li, Na, K, Rb, Cs, Fr (ns¹)

Occurrence:

Na and K are 7^th and 8^th most abundant elements in earth’s crust.

Li= spodumene [LiAl(SiO₃)₂]; 

Na= as NaCl (rock salt), Na₂B₄O₇.10H₂O (borax), Na₂CO₃.10H₂O (washing soda), NaNO₃ (saltpetre), Na₂SO₄; 

K= KCl, K₂SO₄.

Physical properties:

1) atomic size –

a) Alkali metals have largest size among all groups.

Group 1 > group 2 elements.

Reason: atomic size decreases across the period because as atomic number increases, nuclear charge increases, force of attraction between nucleus and outermost electron increases and hence size decreases.

b) Down the group, size increases as number of shells increases.

Order of ionic and atomic radii: Li < Na < K < Rb < Cs.

2) Ionization enthalpy –

a) Alkali metals have low ionization energies.

Reason: as size decreases across the period, force of attraction increases between nucleus and outermost electrons and hence energy required to remove an electron from outermost shell. So, IE increases.

b) Down the group, IE increases as size increases. Li > Na > K > Rb > Cs.

3) Oxidation state = +1. Alkali metals are highly electropositive metals.

4) They are soft metals and can be cut by knife.

5) Metallic character increases down the group.

6) They have high melting points and boiling points as their intermetallic bonds are weak.

7) All alkali metals form ionic bonds except Li (Li forms covalent bonds) and have low densities.

8) Na and K are highly reactive and catch fire easily by reacting with air and moisture. So, they are kept immersed in kerosene oil.

9) Flame colouration – All alkali metals impart colour to flame.

Li = Crimson, Na = Yellow, K = Pale violet, Rb = Violet, Cs = Blue.

Reason: On heating the alkali metal or its salt in a flame, the electrons are excited easily to higher energy levels because of absorption of energy. When these excited electrons return to their ground state, they emit energy which falls in visible region and imparts a characteristic colour to the flame. Different colours appear due to the difference in the amount of energies emitted. Visible region wavelength = 400nm to 750nm.

Note: Be and Mg do not show flame colouration because they have small size and high ionization energy. Hence, it is difficult for them to lose electrons.

10) Photoelectric effect –

It is the phenomenon of emission of electrons from metal surface when electromagnetic radiation of sufficient energy are made to strike against it. All alkali metals exhibit photoelectric effect except Li as Li has very small size and high ionization energy, so it does not release electron easily when exposed to light.

Chemical properties:

1)     Reaction with oxygen – alkali metals form basic oxides with oxygen.                There are three forms of oxides are – oxide (O^2-), peroxide (O₂^2-) and superoxide (O₂^-).

Oxides: 4 Li + O₂ → 2Li₂O

               4 Na + O₂ → 2 Na₂O

               4 K + O₂ → 2K₂O

Peroxides: 2 Na + O₂ → Na₂O₂

                  2 K + O₂ → K₂O₂

Superoxides: K + O₂ → KO₂

Ø  Lithium can form only oxide (Li₂O) and cannot form peroxide and superoxide because of its very small size. Li⁺ has strong positive field (high charge density) around it which attracts the negative charge so strongly that it does not allow oxide anion O^2- to combine with another oxygen atom to from peroxide ion, O₂^2-.

While Na⁺ ion has larger size than Li⁺, so it has weaker positive field around it which allow oxide anion to combine with another oxygen atom to form peroxide ion.

Ø  Superoxides are paramagnetic and hence are coloured due to the presence of unpaired electrons in π^*2p_x² = π^*2p_x¹. Normal oxides are diamagnetic and re colourless.

2)     Reaction with hydrogen – These elements form hydrides.

2 M + H₂ → 2 MH

Such as 2 Na + H₂ → 2 NaH

Order of reactivity of alkali metals towards hydrogen decreases down the group because lattice energy of hydrides decreases as the size of cation increases while going down the group and hence stability of hydrides decreases.

Stability order: LiH > NaH > KH > RbH > CsH.

3)     Reactivity and electrode potential –

Electrode potential is a measure of the tendency of an element to lose electrons in the aqueous solution.

Ø  Alkali metals act as strong reducing agents because they have to lose only one electron from their valence shell and have large atomic size and low ionization energy.

Ø  Down the group, ionization energy decreases so tendency to lose electrons increases and hence, reducing character increases down the group. Order: Na < K < Rb < Cs < Li.

Ø  Lithium is the strongest reducing agent in aqueous solution.

Reason:

Reducing character depends upon ionisation energy and tendency to lose electron from isolated gaseous atom.

Factors affecting: (i) Sublimation: Na(s) → Na(g)

(ii) Ionisation Enthalpy: Na(g) → Na⁺(g) + e^-

(iii) Heat of hydration.

More the tendency of an element to lose electrons, more is the reducing character.

But in case of Lithium, it has smallest size and high heat of hydration enthalpy which compensates the ionisation energy required to remove electron. So, it can easily remove an electron and behaves as a strong reducing agent. Lithium has most negative standard electrode potential. (E⁰ = -3.05V)

4)     Reaction with water – These elements form hydroxides and react vigorously with water.

Ø          2 Na + 2 H₂O → 2 NaOH + H₂,                   2 K + 2 H₂O → 2 KOH + H₂

Ø          Na₂O + H₂O → 2 NaOH

Ø                  K₂O₂ + 2 H₂O → 2 KOH + H₂O₂

Ø                 2 KO₂ + 2 H₂O → 2 KOH + H₂O₂ + O₂.

5)     Reaction with halogens – These elements form halides.

2M + X₂ → 2MX

Reactivity order: Li < Na < K < Rb < Cs & F₂ > Cl₂ > Br₂ > I₂.

6)     Covalent character and polarisation – When a cation approaches an anion, electron cloud of anion is attracted towards cation and gets distorted. This effect is called polarisation.

The power of cation to polarise anion is called polarisation power and tendency of anion to get polarised is called its polarizability.

Greater the polarisation, greater is the covalent character.

By Fajan’s rule – Covalent character increases if:

(i)            The size of cation is small.

Order of covalent character: LiCl > NaCl > KCl > RbCl > CsCl

(ii)          Anion is larger in size.

Order of covalent character: LiI > LiBr > LiCl > LiF

(iii)         Charge of an ion is high.

Order of covalent character: NaCl < MgCl₂ < AlCl₃.

7)     Lattice energy – Amount of energy required to separate one mole of solid ionic compound into its gaseous ions or vice versa.

Greater the lattice energy is, higher is the melting point and boiling point of a compound, greater is the stability of the compound and lesser is the solubility.

Order of stability: LiF > LiCl > LiBr > LiI.

Bond strength of LiF is higher because of small sized cation and anion forming a stable lattice and has higher lattice enthalpy.

8)     Hydration energy – It is the amount of energy released when one mole of gaseous ions combined with water to form hydrated ions.

M⁺ (g) + aq → M⁺ (aq) + Hydration energy

Higher the hydration energy of ions, greater is the solubility of compound in water.

Smaller the size of an ion, more it can be hydrated and greater is its hydration radius and has less ionic mobility.

Hence, Li salts are most hydrated.

Ionic Radius: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Hydrated Radius: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Ionic Mobility: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺.

9)     Solutions in liquid ammonia – All alkali metals dissolve in liquid ammonia giving highly conducting deep blue solutions.

As these solutions contain ammoniated cations and ammoniated electrons, they are good conductors of electricity.

M + (x+y) NH₃ → M⁺(NH₃)_x + e^-(NH₃)_y

When light falls on these electrons, they get excited to higher energy levels by absorbing energy and when they come back to ground state, they impart blue color. In concentrated solution, blue color changes to bronze and becomes diamagnetic.

10)  Salts of oxoacids – Oxoacids are those in which acidic proton is on a hydroxyl group with an oxo group attached to the same atom.

For e.g. H₂SO₄, H₂SO₃, H₂CO₃

Alkali metals form oxoacid salts which are soluble in water and thermally stable.

Stability order: LiCO₃ < Na₂CO₃ < K₂CO₃ < Rb₂CO₃ < Cs₂CO₃.

As Li⁺ is a smaller cation and CO₃^2- is a larger anion. So, Li₂CO₃ has less lattice energy and is stable.

Li₂CO₃ → Li₂O + CO₂

Note: Larger anions = CO₃^2-, HCO₃^-, SO₄^2-, SO₃^2-, I^-, HSO₃^-‑ .

Uses:

Li metal is use to make alloys with Al to make aircraft parts and with Mg to make armour plates. Li is used in electrochemical cells.

Liquid sodium metal is used as a coolant in nuclear reactors.

Potassium is used as an important constituent of fertilizers. KOH is used for absorbing CO₂.

Biological importance of Na & K:

Na⁺ and K⁺ ions develop a sodium-potassium pump which participate in the transmission of nerve signals by developing a potential across cell membranes.

These ions also help in regulating the flow of water, sugars and amino acids across cell membranes. Na⁺ ions are present outside the cells while K⁺ ions are present within the cells.

K⁺ ions activate many enzymes, participate in oxidation of glucose to produce ATP.

Anomalous properties of Lithium: Different properties than other alkali metals,

(i)            Li is much harder and has high melting and boiling point.

(ii)          Li is least reactive but it is the strongest reducing agent among all alkali metals.

(iii)         LiCl is deliquescent and forms LiCl.2H₂O.

(iv)         Li forms normal oxide only and cannot form peroxide & superoxide.

(v)          Li carbonates, hydroxides and nitrates on heating decomposes to give oxide while other metal carbonates & hydroxides do not decompose at all whereas alkali metals nitrates give nitrites on heating.

Li₂CO₃ → Li₂O + CO₂

2LiNO₃ → Li₂O + 2NO₂

NaNO₃ → NaNO₂ + ½ O₂.

Diagonal relationship: Similarities between Li and Mg:

(i)            Both Li and Mg are harder and lighter in their respective groups.

(ii)          Both of them form nitride by direct combination of Li and Mg with N₂ [ Li₃N, Mg₃N₂]

(iii)         Both form normal oxide.

(iv)         LiCl and MgCl₂ are deliquescent and form hydrates. LiCl.2H₂O, MgCl₂.8H₂O.

(v)          Both react slowly with water.

(vi)         Their hydroxides and oxides are less soluble in water.

Note: Learn preparation, physical properties and chemical properties of some important compounds of Na. (Given in NCERT)

Group II: Be, Mg, Ca, Sr, Ba, Ra

General configuration = ns²

General Characteristics:

Physical properties:

(i)            Density – Harder & denser than alkali metals due to smaller size and strong intermetallic bonds.

(ii)          Melting point and boiling point – Higher than alkali metals but no regular trend.

(iii)         Ionization energy and electropositive character – Lower I.E. as compared to p-block elements but it decreases down the group. They are less electropositive than alkali metals. 1^st I.E. is lower than 2^nd I.E.

(iv)         Atomic size and ionic size – They are smaller than alkali metals due to greater nuclear charge.

(v)          Oxidation state – only +2 (high lattice or hydration energy compensates for 2^nd I.E.)

(vi)         Reducing properties – They are strong reducing agents but weaker than alkali metals and reducing character increases down the group.

(vii)        Flame coloration – Be & Mg do not impart colour to the flame as flame energy is insufficient for excitation, the colours of the flame are as follows: Ca – Brick red, Sr – Crimsom, Ba – Apple green, Ra – Crimson.

Chemical Properties:

(i)            Reaction with air and water – They are less reactive then alkali metals. Be & Mg are kinetically inert towards air & water because of the formation of a film of oxide on their surface. Be does not react with steam even at red heat & does not gets oxidized in air below 873K though powdered Be reacts to form BeO & Be₃N₂. The reactivity towards O₂ increases down the group. BeO is amphoteric whereas other oxides are basic. The basic character increases down the group. BeO and MgO are almost insoluble in water due to high lattice energy.

(ii)          Solubility in water – In case of carbonates and sulphates (large anions), stability increase and solubility decreases down the group.

(iii)         Halides – They react directly with halogens to give MX₂.

BeCl₂ : It is prepared indirectly from its oxide.

BeO + C + Cl₂      →  BeCl₂ + CO

Ø  In solid state, structure of BeCl₂ is polymeric with two chlorine atoms are covalently bonded. In vapour phase below 1200K, it exists as dimer and above 1200K, it is a linear monomer.

Ø  Beryllium halides are covalent and all other metal halides are ionic in nature. Ionic character increases down the group.

Ø  Beryllium halides are Lewis acids as the octet of central atom is incomplete.

Ø  Beryllium halides are soluble in organic solvents.

Ø  Fluorides of other alkaline earth metals are insoluble in water due to high lattice energy but chlorides, bromides and iodides are soluble in water.

Ø  Anhydrous halides are hygroscopic and form hydrates such as MgCl₂.6H₂O, CaCl₂.6H₂O, etc.

Ø  Beryllium chloride fumes in moist air due to hydrolysis.

Anomalous behaviour of Be:

It shows anomalous behaviour due to small size, high charge over mass ratio (high polarization power), strong intermetallic bonding, high hydration energy and high ionization energy.

(i)            Beryllium salts are less stable.

(ii)          Be does not react with water.

(iii)         Be is an amphoteric metal.

(iv)         Be does not impart colour to the flame.

(v)          Beryllium carbide react with water to give methane & other carbides give acetylene.

Diagonal Relationship between Be & Al:

(i)            Be & Al are not attacked by acids because of an oxide layer on them.

(ii)          Be & Al form floro complex anions [BeF₄]^2- & [AlF₆]^3- and other alkaline earth metals d not form floro complex.

(iii)         Oxide and hydroxide of Be and Al are amphoteric in nature.

(iv)         BeCl₂ is covalent, polymeric and bridged structure, AlCl₃ is also covalent, bridged and dimer. Both are soluble in organic solvents and are lewis acids.

Important reactions:

(i)            2 BeCl₂ + LiAlH₄ → 2 BeH₂ + LiCl + AlCl₃

(ii)          BeO and Be(OH)₂ are amphoteric.

BeO + H₂O → Be(OH)₂

Be(OH)₂ + 2 OH^- → [Be(OH)_4]^2-

Be(OH)₂ + 2 HCl + 2 H₂O → [Be(OH)₄]Cl₂

(iii)         2 Be(NO₃)₂ → 2 BeO + 4 NO₂ + O₂

2 Mg(NO₃)₂ → 2 MgO + 4 NO₂ + O₂

Biological importance of Mg & Ca:

(i)            Mg – main component of chlorophyll and ATP.

(ii)          Ca – helps in blood coagulation, helps in muscle contraction & relaxation, in nerve transmission, main constituent of bones & teeth.

Uses: Mg is used in making alloys (being lighter in weight).

Mg(OH)₂ is used as an antacid (milk of magnesia).

Note: Learn preparation, physical properties and chemical properties of some important compounds of Ca. (Given in NCERT)