s BLOCK Elements jee

s-Block Elements

The s-block elements

s-Block Elements

Similarities

Highly reactive metals

Strong reducing agents

Form ionic compounds

Fixed oxidation state

Group I : +1

Group II : +2

4Variation in Physical Properties of s-block Elements

1. Atomic Radius and Ionic Radius

2. Ionization Enthalpies

3. Hydration Enthalpies

4. Melting Points

5. Electronegativity

5Atomic and Ionic Radii

The atoms and ions of alkali metals are largest in their

corresponding periods.

Atomic size Li < Na < K < Rb < Cs

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

Atomic volume Li < Na < K < Rb < Cs

Charge density Li > Na > K > Rb > Cs

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6Atomic Radius and Ionic Radius

Group I

element

Atomic radius

(nm)

Group II

element

Atomic radius

(nm)

Li

Na

K

Rb

Cs

0.152

0.186

0.231

0.244

0.262

Be

Mg

Ca

Sr

Ba

0.112

0.160

0.197

0.215

0.217

down the groups the outermost electrons are further away from the nuclei

Group II < Group I ENC from left to right across the periods

7On moving down the groups,

first sharply (e.g. from Li to K)

then slowly (e.g. from K to Fr)

There is a sharp in NC from 19K to

37Rb

Outermost e is drawn closer to the nucleus

The inner d-electrons (of Rb, Cs, Sr, Ba) have poor shielding effect on the outermost electrons transition contraction

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8Ionisation Enthalpy

9Ionization Enthalpy

Both atomic radius and ENC down the groupsAtomic radius is more important

IE down the groups

Group I

element1st IE 2nd IE

Group II

element1st IE 2nd IE 3rd IE

Li

Na

K

Rb

Cs

519

494

418

402

376

7 300

4 560

3 070

2 370

2 420

Be

Mg

Ca

Sr

Ba

900

736

590

548

502

1 760

1 450

1 150

1 060

966

14 800

7 740

4 940

4 120

3 390

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Group I

element1st IE 2nd IE

Group II

element1st IE 2nd IE 3rd IE

Li

Na

K

Rb

Cs

519

494

418

402

376

7 300

4 560

3 070

2 370

2 420

Be

Mg

Ca

Sr

Ba

900

736

590

548

502

1 760

1 450

1 150

1 060

966

14 800

7 740

4 940

4 120

3 390

Ionization Enthalpy

For Group I elements, 2nd IE >> 1st IE because

the 2nd electron is closer to the nucleus and is poorly shielded by other electrons in the same shell which is completely filled.

For Group II elements, 3rd IE >> 2nd IESimilar reasons can be applied

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Variations in the first and second ionization enthalpies of Group I elements

Ionization Enthalpy

12

Variations in the first, second and third ionization enthalpies of Group II elements

Ionization Enthalpy

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Electronegativity

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Electronegativity

Relatively LOW Electronegativity

These metals have more tendency to lose electron rather than to gain an

electron.

The electronegativity values decreases down the group from Li to Cs

Li > Na > K > Rb > Cs

Group I

element

Electronegativity

value

Group II

element

Electronegativity

value

Li

Na

K

Rb

Cs

1.0

0.9

0.8

0.8

0.7

Be

Mg

Ca

Sr

Ba

1.5

1.2

1.0

1.0

0.9

All have low electronegativity => Electropositive

ELECTRONEGATIVITY

EN down the group

EN : Group II > Group I ( greater ENC)

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Hydration

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Hydration enthalpy

Hydration enthalpy (Hhyd) is the amount of energy released when one mole of aqueous ions is formed from its gaseous ions.

M+(g) + aq M+(aq) H = Hhyd

M2+(g) + aq M2+(aq) H = Hhyd

always has a negative value

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Hydration energies

Group I

Alkali metal ions are highly hydrated.

The smaller the ionic size, the higher the degree of hydration.

Primary and secondary shell of hydration

Li ion is very small, it is heavily hydrated.

Li ion is tetrahedrally surrounded by four water molecules using its four sp3

hybrid

Group 2

They have higher hydration energies than Alkali metals due to smaller sizes

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Hydration energies

In aqueous solutions, degree of hydration decreases from Li+ to Cs+ due

to increase in size

Ionic radii of hydrated alkali metal ions also decreases from Li+ to Cs+

Formation of hydrated salts :

Li > Na > K Salts.

Rb and Cs salts are not hydrated

Ionic Mobility : Cs+ > Rb+ > K + > Na + > Li +

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Group I

ion

Hydration

Enthalpy

(kJ mol1)

Group

II ion

Hydration

enthalpy

(kJ mol1)

Li+

Na+

K+

Rb+

Cs+

519

406

322

301

276

Be 2+

Mg2+

Ca2+

Sr2+

Ba2+

2 450

1 920

1 650

1 480

1 360

Group II > Group I Group II ions have higher charge and small

size higher charge density

stronger ion-dipole interaction

Hydration energies

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The melting points of s-block elements depend on the metallic bond strength which in turn depends on

1. charge density of cations

2. number of valence electrons participating in the sea of electrons

3. packing efficiency of the crystal lattices

Melting & Boiling Point

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Group I

element

Melting

Point (C)

Group II

element

Melting

Point (C)

Li

Na

K

Rb

Cs

Fr

180

97.8

63.7

38.9

28.7

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Be

Mg

Ca

Sr

Ba

Ra

1280

650

850

768

714

697

down the groups ionic radii down the groups

charge density interaction between ions and electron sea

Group II > Group I(a) Group II cations have higher charge density

(b) More valence electrons are involved in the sea of electrons(c) Packing efficiency : Group II > Group I

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Group

I

Densit

y (g

cm3)

Group

II

Densit

y (g

cm3)

Li 0.53 Be 1.86

Na 0.97 Mg 1.74

K 0.86 Ca 1.55

Rb 1.53 Sr 2.54

Cs 1.90 Ba 3.59

Fr - Ra -

STRUCTURE

Densities of Li, Na, K

are lesser than that

of water

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Density

Alkali metals have low density.

The reason this is that they have large atomic sizes.

Density gradually increases on moving down the group from Li to

Cs

Anomaly: K is lighter than Na

Li < Na < K < Rb < Cs.

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Effect of light

Alkali metals when irradiated with light emit electrons with ease

due to low ionization enthalpies.

This phenomenon is used in photoelectric cells, particularly

caesium and potassium are used as electrodes in photoelectric

cells.

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Flame Colouration

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Most s-block elements and their compounds give a characteristic flame colour in the flame test

Group I element

Flame colour

Li Crimson

Na Golden yellow

K Lilac

Rb Bluish red

Cs Blue

Flame Colouration

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Most s-block elements and their compounds give a characteristic flame colour in the flame test

Group II element

Flame colour

Be -

Mg -

Ca Brick red

Sr Blood red

Ba Apple green

Beryllium and magnesium atoms are smaller and their electrons being strongly bound to the nucleus are not excited to higher-energy levels.

Flame Colouration

Ca, Sr, Ba

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Mechanism : -

1. In the hotter part of the flame,

2. In the cooler part of the flame,

Na(g) Na(g)*heat

Na(g)* Na(g)cool

[Ne] 3p1 [Ne] 3s1

Ground state

[Ne] 3s1 [Ne] 3p1

+ golden yellow light

Visible region

Flame Colouration

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Mechanism : -

For salts of s-block elements,the metal ions of the salts are first converted to metal atoms

Na+Cl Na(g) + Cl(g)heat

Na(g) Na(g)*heat

Na(g)* Na(g)cool

+ golden yellow light

Na2CO3(s) Na+Cl (more volatile)

Conc. HCl

Flame Colouration

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Complex Formation

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Complex formation

In order to form complex compounds, a metal must possess the

following characteristics.

Small size

High effective nuclear charge

Tendency to accept electrons (i.e., presence of vacant orbitals)

Since alkali metals have none of these characteristics they have little

tendency to form complexes.

Lithium and Beryllium forms certain complexes. (Due to their small sizes)

The complex forming tendency fall markedly down the groups as the

atomic size increases

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Reasons

1. Absence of low-lying vacant d-orbtals to accept lone pairs from ligands.

For Na+, 1s2, 2s2, 2p6, 3s, 3p, 3dHigh-lying relative to 2p

For Fe2+, 1s2, 2s2, 2p6, 3s2, 3p3, 3d6

Low-lying relative to 3p

Complex formation : Weak Tendency

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Reasons

2. s-block cations (M+, M2+) have relatively low charge densities

less polarizing and less able to accept lone pairs from ligands.

Complex formation : Weak Tendency

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Owing to its high charge density, Be2+ can form complexes

Complex formation :

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Electropositive Character

The electropositive character increases down the

group from Li to Cs because ionization enthalpy

decreases down the group

Li > Na > K > Rb > Cs.

Group I (V) Group II (V)

Li -3.04 Be -1.69

Na -2.72 Mg -2.37

K -2.92 Ca -2.87

Rb -2.99 Sr -2.89

Cs -3.02 Ba -2.90

oEMetallic charater (Reactivity)

Group I > Group II

down the groups

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Reducing Property

Powerful reducing agents

Li > Na < K = Rb > Cs (E0)

Reasons

Heat of sublimation

Ionisation enthalpy

Hydration energy

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Reaction with Hydrogen

Group I

2M(s) + H2(g) 2MH(s)300C 500C

Alkali metals react with hydrogen to form ionic hydrides M+H-.

The reaction of alkali metals with hydrogen decreases from Li to Cs


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Group 1 : Hydrides

The order and reactivity with hydrogen

Li > Na > K > Rb > Cs

The ionic character of the bonds in these hydrides Increases from

LiH to CsH

LiH < NaH < KH < RbH < CsH

Stability

LiH > NaH > KH > RbH > CsH

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LAH

Powerful reducing agent

Tetrahedral

Selective reducing agent

Reduces carbonyl compounds to alcohols.

It reacts violently with water, so it is necessary to use absolutely

dry organic solvents

Also reduces several inorganic substances

4LiH + AlCl3 LiAlH4 + 3LiClDry ether

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Sodium tetrahydridoborate (sodium borohydride)

NaBH4

Can be used even in aqueous solutions

Na and K hydrides are useful

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Group II

M(s) + H2(g) MH2(s)

600C 700C

Alkaline earth metals react with hydrogen to form ionic hydrides M2+ (H-)2


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Group 2: Hydrides

Form hydrides of type MH2

Be, Mg Little tendency

Polymeric hydrides (BeH2 )

Three centre two electron bond

BeH2 is covalent

MgH2 is partially ionic

Ca, Ba, Sr ionic hydrides

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Reactions of hydrides

MH(s)

MOH(aq) + H2(g)

MCl(aq) + H2(g)

H (a strong base) tends to react with protonic reagents to release H2

Reactivity down the groups

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Reaction with Air / Oxygen

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Reaction with Air / Oxygen

All alkali metals form more than one type of oxide on burning in air (except lithium)

Group I Elements

All alkaline earth metals react slowly with air to form oxides

On burning in air, they form both oxide and nitride

Group II Elements

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Three types of oxides:

normal oxides

peroxides

superoxides

Reaction with Air / Oxygen : Group 1 Elements

2O

2

1

O2

oxide ion

O22

peroxide ion

2O 2O2

superoxide ion

Abundant supply

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Type of oxide formed depends on

1. supply of oxygen

2. reaction temperature

3. charge density of M+


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Lithium

when it is burnt in air, it forms normaloxide only

C180

4Li(s) + O2(g) 2Li2O(s)lithium oxide


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Sodium

when it is burnt in an abundantsupply of oxygen

forms both the normal oxide and the peroxide

C180

4Na(s) + O2(g) 2Na2O(s)sodium oxide

C300

2Na2O(s) + O2(g) 2Na2O2(s)sodium peroxideexcess


52

Potassium, rubidium and caesium

form All three types of oxides when burnt in sufficient supply of oxygen


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Group I

elementNormal oxide Peroxide Superoxide

Li

Na

K

Rb

Cs

Li2O

Na2O

K2O

Rb2O

Cs2O

Na2O2

K2O2

Rb2O2

Cs2O2

KO2

RbO2

CsO2

Cations with high charge densities (Li+ or Na+) tend to polarize the large electron clouds of peroxide ions and/or superoxide ions

Making them decompose to give oxide ions


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The electron cloud of the superoxide ion is greatly distorted by the small lithium ion


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Group I

elementNormal oxide Peroxide Superoxide

Li

Na

K

Rb

Cs

Li2O

Na2O

K2O

Rb2O

Cs2O

Na2O2

K2O2

Rb2O2

Cs2O2

KO2

RbO2

CsO2

Super oxides are generally bright coloured

They exhibit paramagnetic character due to unpaired electron


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KO2 used as oxygen generators and CO2 scrubbers in spacecrafts and submarines

4KO2 + 2H2O 4KOH + 3O2

2KOH + CO2 K2CO3 + H2O


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Group II

elementNormal oxide Peroxide

Supero

xide

Be

Mg

Ca

Sr

Ba

BeO

MgO

CaO

SrO

BaO

-

-

All these oxides are basic in nature (except beryllium oxide which is amphoteric)


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Solubility

G0 = H0 - TS0

General rule:

Compounds that contain widely differing radii are soluble in water

Difference in size favours solubility (>80pm)

Thermodynamics of dissolution

Entropy favours dissolution

Hydration energy of a smaller ion is larger

LH = 1 / (r+

+ r-) and HydH = (1 / r

+) + (1 / r

-)

Ion size assymmetry results in exothermic dissolution

If both are small, both LH and HydH may be large, but enthalpy of

dissolution may not be very exothermic

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Solubility

The solubility of compounds increases with increase in ionic size of

metal

Fluorides, oxides, hydroxides

The solubility of compounds decreases with increase in ionic size of

metal

Carbonates, sulphates, nitrates, halides (except fluorides)

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Two processes are

1. Breakdown of the ionic lattice

2. Hydration

Processes involved in Dissolution and their Energetics

61

NaCl(s) Na+(aq) + Cl

-(aq)

Na+(g) + Cl

-(g)

Hsolution

olattice

ohydration

osolution HHH

= (-772 +776) kJ mol1

= +4 kJ mol1

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osolution

osolution

osolution STHG

If , we expect the solids to dissolve in water

0Hosolution

Solubility as becomes more ve (less +ve)osolutionH

Solids (e.g. NaCl) with small +ve valuesare also soluble in water if the dissolution involves an increase in the entropy of the system.

osolutionH

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osolution

osolution

osolution STHG

0Gosolution Spontaneous dissolution

osolutionST is always positive

osolutionHDissolution with slightly positive

can be spontaneous

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Trends and Interpretations

1. The solubility of Group(II) sulphate decreases down the group

On moving down the group, cationic radius(r+)

both and become less -veoLH

ohydrationH

However, less rapidly than oLH

ohydrationH

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rr

1H

24SO

oL

rr 24SO

constant

olattice

ohydration

osolution HHH

less ve down the group

+ve constantless ve down the group

Solubility down the group

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rr

1H

24SO

oL

rr 24SO

constant

olattice

ohydration

osolution HHH

more rapidly down the group

less rapidly down the group

less ve down the group


(-ve) (+ve)

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2. The solubility of Group(II) hydroxides increases down the group

On moving down the group, cationic radius(r+)

both and become less -veoLH

ohydrationH

However, more rapidly than oLH

ohydrationH

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olattice

ohydration

osolution HHH

less rapidly down the group

more rapidly down the group

more ve down the group


(-ve)(+ve)

less +ve down the group

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General Rules

For s-block compounds with small anions (e.g. OH, F),

solubility in water down the group

For s-block compounds with large anions (e.g. SO42, CO3

2-),

solubility in water down the group

For s-block compounds with medium size anions (e.g. Br),

solubility in water exhibits irregular pattern down the group

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Group II compounds with doubly-charged anions (MX) are less soluble than those with singly-charged anions (MY2)

Reasons :

1. HL of MX > HL of MY2

2. HL is the major factor affecting solubility

Hsolution of MX is more positive

Solubility : MX < MY2

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Solubility : Group I > Group II

Reasons :

For a given anions, both HL and Hhydration become more ve from Group I to Group II

However, HL is the major factor affecting solubility

Hsolution : Group I is less positve than Group II

Solubility : Group I > Group II

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Thermal Stability

G0 = H0 - TS0

The G0 for the decomposition of a solid becomes negative when

TS0 > H0

H0 depends on (example carbonates)

= Enthalpy of decomposition + (Lattice Enthalpy of Carbonate - Lattice

Enthalpy of Oxide)

Enthalpy of decomposition is generally large and positive

Metals having small cations, increases the lattice enthalpy of oxide

more than that of the carbonate / sulphate / hydroxide / peroxide

Therefore, Lattice enthalpy plays an important role in deciding the

stability.

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Thermal decomposition reactions

Metal carbonates

M2CO3(s) M2O(s) + CO2heat

MCO3(s) MO(s) + CO2heat

Metal hydroxides

2MOH(s) M2O(s) + H2O(g)heat

M(OH)2(s) MO(s) + H2Oheat

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Relative thermal stability can be measured in two ways

A higher decomposition temperature

a greater thermal stability

C100

BeCO3(s) BeO(s) + CO2(g)

MgCO3(s) MgO(s) + CO2(g) C540

CaCO3(s) CaO(s) + CO2(g) C900

SrCO3(s) SrO(s) + CO2(g) C1290

BaCO3(s) BaO(s) + CO2(g) C1360

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Relative thermal stability can be measured in two waysBy comparing the standard enthalpy changes of thermal decomposition

reactions

A more positive H value a thermally more stable compound

M(OH)2(s) MO(s) + H2O(g) H > 0

Be(OH)2(s) BeO(s) + H2O(g)H = +54 kJ mol1

Mg(OH)2(s) MgO(s) + H2O(g)H = +81 kJ mol1

Ca(OH)2(s) CaO(s) + H2O(g)H = +109 kJ mol1

Sr(OH)2(s) SrO(s) + H2O(g)H = +127 kJ mol1

Ba(OH)2(s) BaO(s) + H2O(g)H = +146 kJ mol1

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Factors affecting thermal stability

1. Polarizing power of cation

2. Polarizability of polyatomic anion

3. Lattice enthalpy of metal oxide produced

77

Interpretation of trends in thermal stability of carbonates and hydroxides

1. Group I > Group II

(a) M2+ ions have higher charge densities than M+ ions

M2+ ions are more polarizing than M+ ions

Can polarize more the electron cloud of polyatomic anions

Polarizability as the size of anion

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1. Group I > Group II

(b) M2+ ions have higher charge densities than M+ ions

Lattice enthalpy : MO > M2O

Energetic stability : MO > M2O

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CaCO3(s) CaO(s) + CO2(g)heat

Na2CO3(s) Na2O(s) + CO2(g)

more favourable

less favourable

heat

more stable

less stable

Thermal stability of carbonates : -

Group I > Group II

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2. Thermal stability down the groups

size of cations down the groups

(a) charge density/polarizing power of cation down the groups

(b) lattice enthalpies of MO/M2O down the groups

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MgCO3(s) MgO(s) + CO2(g)heat

more favourable

more stable

BaCO3(s) BaO(s) + CO2(g)heat

less favourable

less stable

more polarized

less polarized

Thermal stability of carbonates

down the groups

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Effect of sizes of the cations on thermal stability of the carbonates and hydroxides of both Groups I and II metals

83

Reaction with Water

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Action of Water

Both alkali and alkaline metals react with water

Respective Hydroxides and Hydrogen gas are formed

Reactivity increases down the group

Type : Slow to explosive reactions

Na K

Reactions with water or steam

Group I

2M(s) + H2O(l) 2MOH(aq) + H2(g)heat

Group II

M(s) + 2H2O(l) M(OH)2(aq) + H2(g)heat

Mg reacts with steam but not cold water

Be has no reaction with either water or steam

Mg(s) + H2O(g) MgO(s) + H2(g)heat

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Reaction with ammonia

Exhibited both by Group I and II metals

All show Blue colour

Ammoniated electron is present in these solutions, as the

electron is solvated by ammonia

Intensity of blue color increases with metal concentrations

High electrical conductivity

This solution show Magnetic properties

Reducing property of solution of metal in ammonia (selective

reducing action in organic chemistry)

These solution scan be used to prepare any desired oxide, by

passing calculated quantities of oxygen gas through the solutions

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Hydroxides

88

Group 1 Hydroxides of type MOH

These hydroxides are Strong bases

Basic strength / basic character / solubility in water / thermal

stability

LiOH < NaOH < KOH < RbOH < CsOH

LiOH decomposes on heating to give water and Li2O

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Group 2: Hydroxides of type M(OH)2

All group 2 metals form hydroxides

Reaction of oxides with water gives hydroxides

Be(OH)2 Mg(OH)2 Ca(OH)2 , Sr(OH)2 , Ba(OH)2

Amphoteric Weakly Basic Strongly Basic

Weaker bases than alkali metal hydroxides

Higher IE, smaller ionic size, higher charge on metal ion.

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Group 2: Hydroxides

The solubility of the hydroxides in water increases with

increase in atomic number of the cation.

Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba (OH)2

insoluble insoluble sp. soluble soluble soluble

The solubility of hydroxides depend mainly on two facts.

The lattice energy required to dissociate the components of hydroxide. This decreases from beryllium to barium.

The hydration energy of cation M2+. This decreases from beryllium to barium as the size of cation increases.

Both lattice and hydration energies decrease down the group, the decrease

in lattice energy is more rapid than the hydration energy and so their

solubility increases on descending the group.

CompoundsSolubility / mol per 100 of

water

Mg(OH)2 0.02 103

Ca(OH)2 1.5 103

Sr(OH)2 3.4 103

Ba(OH)2 15 103

down the group


water

MgSO4 1800 104

CaSO4 11 104

SrSO4 0.71 104

BaSO4 0.009 104

down the group


water

Mg(OH)2 0.02 103

Ca(OH)2 1.5 103

Sr(OH)2 3.4 103

Ba(OH)2 15 103


water

MgSO4 1800 104

CaSO4 11 104

SrSO4 0.71 104

BaSO4 0.009 104

Size and/or charge of the anion

Polarizability of anion

Covalent character

Solubility in water

In general,

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Group I: Carbonates

Type M2CO3

Solubility in Water:

Increases as the size (atomic number) of cation increases.

Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3

Low --------------High-------------------------------

Thermal Stability

Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3

Low --------------High-------------------------------

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Group II: Carbonates

All form carbonates of the type MCO3

Solubility in Water:

Insoluble in neutral medium, soluble in acidic medium

Solubility decreases down the group

BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3

Carbonates are more soluble in a solution containing CO2 Bicarbonates

All carbonate solutions undergo the above reaction

Bicarbonates cannot be obtained in solid form but are known in solution state

only.

Na, K, Rb, Cs bicarbonates are the only ones that can be obtained in solid

state.

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Group II: Carbonates

Thermal Stability Increases down the group

BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3

BeCO3 must be stored under CO2

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Reaction with Nitrogen

97

Alkali Metal : Reaction with Nitrogen family

Lithium forms Nitrides (exceptions w.r.t alkali metal reactions)

Other metals form Azides (MN3)

They form binary compounds with other family members of N

The binary compounds undergo hydrolysis in water to form

ammmonia, phosphine, asine, stibine etc

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Alkaline Earths : Reaction with nitrogen family

All metals form nitrides M3N2

Ease of formation of nitrides decreases down the group

These nitrides are stable up to 10000C

Get hydrolysed in water to give ammonia

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Halides

100

Group 1 : Halides

MX

Ionic compounds , high lattice energies,

Stability

The order of enthalpy of formation of a metal halide is

Fluoride > Chloride > Bromide > Iodide

Fluorides are highly stable

2M(s) + Cl2(g) 2MCl(s)heat

101

Group 1: Halides

Trends in melting and boiling points of halides:

For a given alkali metal, the melting points and boiling points:

Fluoride > Chloride > Bromide > Iodide

For a given halogen, the melting and boiling points

Lithium < Sodium > Potassium > Rubidium > Caesium Due to covalent character of Li compound

102

Halides : Ionic Character

The order of ionic character is

LiX < NaX < KX < RbX < CsX

MF > MCI > MBr > MI

(same metal, different halogen)

103

Group 2: Halides

MX2

When crystallized from solutions they form hydrated salts

Anhydrous CaCl2, SrCl2 and BaCl2 can be prepared by heating the

hydrated salts.

M(s) + Cl2(g) MCl2(s)heat

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Group 2: Halides

Alkaline earth metals combine with halogen on heating

to form MX2 type salts.

Be Halides are covalent

Other halides are ionic

Ionic Character

BeX2 < MgX2 < CaX2 < SrX2 < BaX2 MI2 < MBr2 < MCl2 < MF2

105

A) At high temperatures, BeCl2 occurs as a gaseous molecule with only four

electrons around Be.

B) In the solid state, BeCl2 occurs in long chains with each Cl bridging two

Be atoms, which gives each Be an octet.

Structure of BeCl2 molecules

106

Group 2: Halides

Except BeCl2, all other halides are hygroscopic

Extent of hydration decreases down the group

Be and Mg halides hydrolyse on heating

Ca, Sr, Ba halides get dehydrated on heating

Calcium chloride has a strong affinity for water

Solubility order

Fluorides are readily soluble

BeF2 > MgF2 > CaF2 < SrF2 < BaF2 BeX2 > MgX2 > CaX2 > SrX2 > BaX2 MF2 < MCl2 < MBr2 < MI2

Saahil Jain

107

Reactions of chlorides

No significant reactions with water, acids or alkalis

Group I

Group II

Do not undergo significant hydrolysis except BeCl2 and MgCl2

BeCl2(aq) + 2H2O(l) Be(OH)2(aq) + 2HCl(aq)

MgCl2(aq) + H2O(l) Mg(OH)Cl(aq) + HCl(aq)

Basic salt

More favoured in alkaline solutions

108

Group II Sulfates

Obtained by action of dil sulfuring acid on

Metal

Metal oxide

Metal hydroxide

Carbonate

Sulfates of Be, Mg, Ca crystalise as Hydrated salts

BeSO4 . 4H2O MgSO4 . 7H2O CaSO4 . 2H2O

Sulfates of Sr and Ba crystallise without water of crystallisation

109

Group II Sulfates

Solubility in water

BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4

fairly soluble completely insoluble

Thermal stability increases down the group

BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO4

110

Nitrates

Group 1 : All form nitrates of type MNO3

Group 2 : All form nitrates of type M(NO)2

All are ionic

All are soluble in water

111

General Reactions of Alkali Metals

112

Important Reactions of the Alkaline Earth Metals - I

The metals reduce O2 to form the oxide:

Barium also forms the peroxide BaO (s).

The Metals of higher atomic weight reduce water to form hydrogen gas:

Be and Mg form an adherent oxide coating that allows only

slight reaction.

2 M(s) + O2 (g) MO(s)

M(s) + 2 H2O(l) M(OH)2 (aq) + H2 (g)M = Ca, Sr and Ba

113

Important Reactions of the Alkaline Earth Metals - I

The metals reduce halogens to form ionic halides:

Most of the metals reduce hydrogen to form ionic hydrides.

M(s) + X2(-) MX2 (s) X = F, Cl, Br, I

M(s) + H2 (g) MH2 (s) all except Be

114

Important Reactions of the Alkaline Earth Metals - II

Most of the metals reduce nitrogen to form ionic nitrides:

Except for amphoteric BeO, the oxides are basic:

All carbonates undergo thermal decomposition to the oxide:

MCO3 (s) MO + CO2

This reaction is used to produce CaO (lime) in huge amounts from naturally occurring limestone, and was the reaction used to generate carbon dioxide to smother the graphite fire in the Chernobyl reactor.

3 M(s) + N2 (g) M3N2 (s) all except Be

MO(s) + H2O(l) M(OH)2 (aq)

115

General Reactions of Alkaline Earth Metals

116

Diagonal relationship

117

ReactionOther Group I

elementsLithium Magnesium

Combination with O2Peroxides and superoxides

Li2O (normal oxide) MgO (normal oxide)

Combination with N2 No reaction Li3N Mg3N2

Action of heat on carbonate

No reaction (thermally stable)

Decomposes to give Li2O and CO2

Decomposes to give MgO and CO2

Action of heat on hydroxide

No reaction (thermally stable)

Decomposes to give Li2O and H2O

Decomposes to give MgO and H2O

Action of heat on nitrate

Decomposes to give MNO2 and O2

Decomposes to give Li2O, NO2 and O2

Decomposes to give MgO, NO2 and O2

Hydrogen carbonates Exist as solids Only exist in solution

Solubility of salts in water

Most salts are more soluble than those of

Li, Mg.

Fluoride, hydroxide, carbonate, phosphate, ethanedioate are sparingly soluble.

Solubility of salts in organic solvents.

Halides only slightly soluble in organic

solvents

Halides (with covalent character) dissolve in organic solvents

119

Similarities of Be and Al

Be and Al have the same electronegativity (Be I.0 and AI 1.5) and their

charge/radius ratios are very

The standard oxidation potentials of both Be and are of nearly the same order (Be

= 1.97V; Al= l.7V)

Since the polarizing power of both Be and Al are nearly the same, the covalent

character of their compounds also similar.

Both Be and Al are rendered passive on treatment with conc. HNO3.

Unlike alkaline earth metals Be does not get readily attacked by dry air. (like Al)

Both Be and Al reacts very slowly with dilute mineral acids due to the presence of

oxide layer.

Both Be and Al react with alkalis liberating H2.

Both Be and AI form carbides which on hydrolrolysis liberate methane.

Both form nitrides when heated in nitrogen which give ammonia by the reaction

with water.

Both form oxides which are amphoteric.

Halides of both Be and AI contain halogen bridge bonds

122

Compounds of Na

123

Na2CO3 Solvay Process

124

Ca

rb

on

atin

gT

ow

er

126

Sodium Carbonate

127

Caustic Soda NaOH - Nelson Cell

128

NaOH by Castner Kellner Cell

129

NaOH by Castner Kellner Cell

Castner-Kellner method also known as Mercury Cathode Method.

In this method, the electrolytic cell contains three compartments.(i) Mercury in the outer compartment acts as a cathode while in middle

compartment acts as an anode due to induction.

(ii) Graphite rods in the outer compartments acts as anode while the iron rods in the middle compartment acts as a cathode.

Sodium liberated at mercury cathode in the out compartments dissolve in mercury forming sodium amalgam which moves into middle compartment

where it react with water at cathode forming NaOH, H2 and Hg. Cl2 gas is liberated at graphite anodes in the outer compartment.

130

Sodium Hydroxide

131

Sodium Hydroxide

132

Sodium Hydroxide

Sodium Beryllate

Sodium Aluminate

Sodium Stannite

Sodium Plumbite

Sodium Zincate

133

Sodium Hydroxide

134

Sodium Sulphate - Salt Cake anhydrous Na2SO4- Glaubers Salt Na2SO4 .10 H2O

Preparation

The salt cake (anhydrous sodium sulphate) is dissolved in water and the solution

is subjected to crystallization.

Above 32 C the anhydrous salt separates.

Below 32 C, the decahydrate salt crystallises out from the aqueous solution.

Saturated solution of the decahydrate, on cooling below 12 C, gives crystals of

heptahydrate.

Properties

Uses: It is used in textile industry, medicines as purgative, manufacture

of glass plates and sodium salts.

344223

4422

344223

2)(

2

2)(

NaNOSrSOSONaNOSr

NaClBaSOSONaBaCl

NaNOPbSOSONaNOPb

135

Sodium Bicarbonate, Baking Soda, Na2HCO3 Preparation

By passing CO2 through Sodium Carbonate solution but industrially it is

manufactured by Solvay's process.

Properties

It is sparingly soluble in Water

Solution is alkaline in nature

Uses

On heating, it decomposes to give sodium carbonate.

The metal salts which gives basic metal carbonate with sodium carbonate

gives normal carbonates.

Sodium bicarbonate is used, as an antacid in medicine, in dry fire

extinguishers, in baking powders and as mild antiseptic for skin infections.

136

Compounds of Alkaline Earth Metals

137

Magneisum Oxide - MgO Magnesia

Preparation:

1. Calcination of Magnesite (MgCO3)

MgCO3 MgO + CO2

2. Heating Mg(NO3)2 or Mg(OH)2

Mg(NO3)2 MgO + 4 NO2 + O2

Mg(OH)2 MgO + H2O

Properties:

Light infusible white solid

They have high MP 3073K

Used as refractory material due to the above property

138

Magneisum Oxide - MgO Magnesia

Chemical properties

Hydrolyses in water to form insoluble Mg(OH)2

Being basic, reacts with acids to form respective salts

Gives Mg on reduction with Carbon at high tempertaures

MgO + C Mg + CO

Uses:

Sorels cement used in Dentistry MgCl2. 5MgO.xH2O

As an antacid

As an insulator when mixed with asbestos

22 )OH(Mg2OHMgO2

OHMgClHCl2MgO 22

COMgCC3MgO 2

139

Magnesium Hydroxide Mg(OH)2

Preparation:

1. By the hydrolysis of MgO

2. By treating MgCl2 with Ca(OH)2

MgCl2 + Ca(OH)2Mg(OH)2 + CaCl2

White powdery substance

Sparingly soluble in water

Used as an Antacid under the name Milk of Magnesia

140

Magnesium Carbonate MgCO3

Preparation:

Hot Magnesium sulfate with sodium bicarbonate

Basic magnesium carbonate

(Basic Magnesium Carbonate / Magnesia alva)

A solution containing 12% MgCO3 per 100 cc of water containing dissolved

CO2 in called Fluid Magnesia

2242334 2 COOHSONaMgCONaHCOMgSO

24223324 )(.2

COSONaOHMgMgCOCONaMgSOOH

232223 )(3)(. HCOMgOHCOOHMgMgCO

22323)( COOHMgCOHCOMg

141

Magnesium Sulphate MgSO4.7 H2O Epsom Salt

Preparation

From Magnesite : heating with dil. Sulfuric acid

From Dolomite

From Keiserite (commercial method). Boil with water and cool

Properties

Colourless efflorescent solid

OHCOMgSOSOHMgCO 224423

OHCOCaSOMgSOSOHCaCOMgCO 2244423.3 222

OHMgSOOHOHMgSO 2.4224 7.

4C200

24C150

24C30

24 MgSOOH.MgSOOH6.MgSOOH7.MgSO

142

Magnesium Sulphate MgSO4.7 H2O Epsom Salt

On heating, it decomposes

On heating with Carbon, it gets reduced

It forms double salts with alkali metal sulfates

232C250

4 OSO2SO2MgO4MgSO4

224 COSO2MgO2CMgSO2

OH6.MgSO.SOK 2442

143

Calcium Oxide / Quick Lime / Burnt Lime

Preparation: Decomposition of Limestone

Reacts with water with a hissing noise to form Slaked lime Ca (OH)2

(Rxn is known as Slaking of lime) H = -15 kcal/mol

Milk of lime : paste of lime in water

Lime water : Clear filtrate

Limelight in oxy hydrogen flame

22 )OH(CaOHCaO

2C900

3 COCaOCaCO

144

CaO gives Calcium Silicate with silica and Calcium Phosphate with P4O10

Forms Calcium Carbide on heating with carbon (2000 deg C)

Calcium Carbide + water gives Calcium Cyanamide

Calcium Cyanamide + C = Nitrolim - a Fertiliser

24352

32

)(226 POCaOPCaO

CaSiOSiOCaO

COCaCC3CaO 2C2000

CCaCNNCaCcyanamideCalcium

222C1000


145

Uses

It is used as a drying agent.

It is used in the manufacture of bleaching powder.

It is used in the manufacture of calcium. carbide, cement, glass, lime

mortar, etc.

It is used in the purification of sugar.


146

Calcium Hydroxide / Slaked Lime / Milk of Lime

Preparation:

Slaking of lime

Properties

White amorphous solid

Sparingly soluble in water

On heating, loses water molecule to form Lime CaO

Action of CO2

Similar reaction with SO2 gas is seen when Calcium bisulphite is

formed

OHCaCOCO)OH(Ca 2322

lelubSolelubInso

)HCO(CaCOOHCaCO 23223

147

Calcium Hydroxide / Slaked Lime / Milk of Lime

Reaction with Ammonia

Reaction with Chlorine

Bleaching powder is a calcium salt of hypochlorous acid (HOCl)

Ca(OCl)2

Uses: It is used

1. for absorbing acid gases.

2. in the manufacture of bleaching powder and caustic soda.

3. in the production of lime mortar for construction of buildings, whitewashing buildings

4. in glass making, tanning industry and for purification of sugar

5. for the preparation of NH3 from NH4Cl in Solvay process

6. as lime water in laboratories

OHNH2CaClClNH2)OH(Ca 232Heat

42

Ca(OCl)2.Ca(OH)

2.CaCl

2. H

2O + H

2O3Ca(OH)2 + 2Cl2

below 35oC

slaked lime bleaching powder

2

148

Gypsum CaSO4.2H2O Preparation

Properties

White crystalline solid

Solubility decreases on increase in temperature

Action of heat Calcium sulphate hemihydrate (plaster of paris) is formed

HCl2CaSOSOHCaCl 4422

NaCl2CaSOSONaCaCl 4422

OH3OH.)CaSO(]OH2.CaSO[2 2224C120

24

OHCaSO2]OH.)CaSO[( 24C200

224

22heatedStrongly

4 OSO2CaO2CaSO2

Plaster of Paris

Dead burnt plaster

149

Calcium Carbonate, CaCO3

Naturally found as limestone, marble, chalk

Preparation

White fluffy powder insoluble in water. But dissolves in water in the

presence of carbon-di-oxide to form calcium bicarbonate

OHCaCOCO)OH(Ca 2322

NaCl2CaCOCONaCaCl 3322

23223 )HCO(CaCOOHCaCO

150

Mortar

It is also known as lime mortar.

It is an intimate mixture of 1 part of slaked lime, 3 parts of sand and

water made into paste.

This is used to bind the bricks firmly.

Setting of mortar involves the following steps.

(i) Mortar loses water on account of evaporation.

(ii) Carbon dioxide is absorbed from the air converting into calcium

carbonate which acts as a binding material.

(iii) Slaked lime reacts with silica forming calcium silicate which gives

hardness.

151

Cement

The name portland cement was given to it by Joseph Aspidin (a mason!)

because when it is mixed with sand and water it hardens like the lime

stone querried at Portland in England.

Composition :

CaO 50 to 60 %; SiO2: 20 to 25%; Al2O3 : 5 to 10 %; MgO :2 to 3%; Fe2O3 I

to 2% and SO3 1 to 2%.

If lime is excess the cement cracks during setting but if it is less the

cement will be weak.

Excess of Al2O3 will make cement quick drying

The raw materials for the manufacture of cement are limestone and

alumino silicates (clay, sand and shales). When the powdered raw

materials are heated in a rotary kiln, sintered clinker will be obtained.

The setting of cement by mixing with water is due to hydration of the

molecules and their rearrangement.

152

S - Block Metals in Biological Systems

153

Biological functions of Sodium and Potassium ions

Sodium and potassium are the most common cations in biological fluids.

Sodium ion is the major cation of extracellular fluids of animals and in blood plasma,

including human beings which is known to activate certain enzymes in the animal

body

These ions participate in the transmission of nerve signals.

They also regulate flow of water across cell membranes and in transport of sugars,

amino acids into the cells.

Potassium ions are the most abundant cations within cell fluids, where they activate

many enzymes that participate in oxidation of glucose to produce adenosine

triphosphate (ATP).

A typical 70 kg adult contains about 90 g of Na+ ions and 170 g of K+ ions.

The daily requirement of sodium and potassium is about 2 g each.

154

Biological functions of Magnesium and Calcium

Magnesium is an important constituent of chlorophyll.

Mg2+ and Ca2+ ions are also responsible for the transmission of electrical

impulses along the nerve fibre and the contraction of muscles

Calcium ions are essential for the formation of bones and teeth

It also plays important roles in maintaining rhythm of heart, clotting of blood,

neuromuscular function, interneuronal transmission, cell membrane integrity,

etc.

The calcium concentration in plasma is regulated at about 100 mg L-1. It is

maintained by two hormones, calcitonin and parathyroid hormone.

The substance present in bones is continuously solubilized and redeposited to

the extent of 400 mg per day in man.

All this calcium passes through the plasma.

The END

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