The s-Block Elements 40.1Characteristic Properties of the s-Block Elements 40.2Variation in...

The The ss-Block Elements-Block Elements

40.140.1 Characteristic Properties of the Characteristic Properties of the ss-Block -Block

ElementsElements

40.240.2 Variation in Properties of the Variation in Properties of the ss-Block -Block ElementsElements

40.340.3 Variation in Properties of the Variation in Properties of the CompCompounds of the ounds of the ss-Block Elements-Block Elements

4040


• Elements of Groups IA* (the alkali metals) and IIA* (the alkaline earth metals)

constitute the s-block elements

their outermost shell electrons are in the s orbital

*Note: In the following, Groups IA and IIA are abbreviated as Groups I and II respectiv

ely.


• The two groups of elements have many similarities

highly reactive metals

strong reducing agents

form ionic compounds with fixed oxidation states of +1 for Group I

elements and +2 for Group II elements

The s-block elements

Group I elementsGroup I elements

• Lithium


• Sodium


• Potassium


• Rubidium


• Francium


• Beryllium


• Magnesium


• Calcium


• Strontium


• Barium


• Radium

40.140.1Characteristic PrCharacteristic Pr

operties of theoperties of thes-Block Elementss-Block Elements

40.1 Characteristic Properties of the s-Block Elements (SB p.38)

Some characteristic properties of Group I elements

Group I element

Atomic number

Electronic configuration

Electronegativity value

Oxidation state in compounds

Li

Na

K

Rb

Cs

Fr

3

11

19

37

55

87

[He] 2s1

[Ne] 3s1

[Ar] 4s1

[Kr] 5s1

[Xe] 6s1

[Rn] 7s1

1.0

0.9

0.8

0.8

0.7

–

+1

+1

+1

+1

+1

–


Some characteristic properties of Group I elements

Group I element

Oxide formedHydroxide

formedFlame colour

Li

Na

K

Rb

Cs

Fr

Li2O

Na2O, Na2O2

K2O, K2O2, KO2

Rb2O, Rb2O2, RbO2

Cs2O, Cs2O2, CsO2

–

LiOH

NaOH

KOH

RbOH

CsOH

–

deep red

yellow

lilac

bluish red

blue

–


Some characteristic properties of Group II elements

Group II element

Atomic number

Electronic configuration


Oxidation state in compounds

Be

Mg

Ca

Sr

Ba

Ra

4

12

20

38

56

88

[He] 2s2

[Ne] 3s2

[Ar] 4s2

[Kr] 5s2

[Xe] 6s2

[Rn] 7s2

1.5

1.2

1.0

1.0

0.9

–

+2

+2

+2

+2

+2

–


Some characteristic properties of Group II elements

Group II element

Oxide formedHydroxide

formedFlame colour

Be

Mg

Ca

Sr

Ba

Ra

BeO

MgO

CaO

SrO, SrO2

BaO, BaO2

–

Be(OH)2

Mg(OH)2

Ca(OH)2

Sr(OH)2

Ba(OH)2

–

–

–

bluish red

blood-red or crimson

blue

–

Metallic CharacterMetallic Character

• All Group I elements

silvery solids and tarnish rapidly in air at room temperature and pressure

stored under paraffin oil or in vacuum-sealed ampoules

(to prevent contact with oxygen and water vapour in air)




weak metallic bonds

only one valence electron per atom delocalized into the electron s

ea for the formation of metallic bonds

soft and can be cut with a knife easily




low melting points and boiling points



sodium

Sodium is stored under paraffin oil


Caesium and rubidium are stored in vacuum-sealed ampoules

caesium rubidium



body-centred cubic structures

much empty space

comparatively low densities



Some information about Group I elements

Group I element

Atomic radius (nm)

Ionic radius (nm)

Metallic structure

Melting point(C)

Lithium

Sodium

Potassium

Rubidium

Caesium

Francium

0.152

0.186

0.231

0.244

0.262

0.270

0.060

0.095

0.133

0.148

0.169

0.176

b

b

b

b

b

–

180.5

97.8

63.7

39.1

28.4

27


Some information about Group I elements

Group I element

Boiling point(C)

Density at 20 C (g cm–3)

Abundance on Earth (%)

Lithium

Sodium

Potassium

Rubidium

Caesium

Francium

1 330

890

774

688

690

680

0.53

0.97

0.86

1.53

1.87

–

0.002 0

2.36

2.09

0.009 0

0.000 10

Trace


• All Group II elements

greyish solids at room temperature and pressure

also be cut with a knife, but harder than the alkali metals




two valence electrons per atom

smaller atomic sizes

metallic bonds are stronger




the melting points and boiling points are higher than those of Group I

elements

show different metallic structures



• Beryllium and magnesium

hexagonal close-packed structures



• Calcium and strontium

face-centred cubic structures



• Barium

body-centred cubic structure



Some information about Group II elements

Group II element

Atomic radius (nm)

Ionic radius (nm)

Metallic structure

Melting point(C)

Beryllium

Magnesium

Calcium

Strontium

Barium

Radium

0.112

0.160

0.197

0.215

0.217

0.220

0.031

0.065

0.099

0.113

0.135

0.140

h

h

f

f

b

–

1 278

648.8

839

769

729

697


Some information about Group II elements

Group II element

Boiling point(C)

Density at 20 C (g cm–3)

Abundance on Earth (%)

Beryllium

Magnesium

Calcium

Strontium

Barium

Radium

2 477

1 100

1 480

1 380

1 640

1 140

1.85

1.75

1.55

2.54

3.60

5.0

0.000 28

2.33

4.15

0.038

0.042

Trace

Low ElectronegativityLow Electronegativity

• All s-block elements

low electronegativity values




electropositive elements

their atoms have a relatively high tendency to lose their outermost s

hell electrons



• The outermost shell electrons

effectively shielded from the nucleus by the fully-filled inner electron sh

ells

the outermost shell electrons are only loosely held by the nucleus



• Going down both Groups I and II

the elements become more electropositive

the atoms tend to lose electrons more readily

the outermost shell electrons are much further away from the

nucleus



• Group II elements

more electronegative than the Group I elements

the increase in effective nuclear charge

the attractive force between the nucleus and the outermost shell electrons becomes stronger



Electronegativity values of Groups I and II elementsGroup I element


Group II element


Li

Na

K

Rb

Cs

Fr

1.0

0.9

0.8

0.8

0.7

–

Be

Mg

Ca

Sr

Ba

Ra

1.5

1.2

1.0

1.0

0.9

–

Formation of Basic OxidesFormation of Basic Oxides


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

1. Group I Elements1. Group I Elements


• Three types of oxides:

normal oxides

peroxides

superoxides


• They are all ionic

• They can be related as follows:



2O21

2OO2–

oxide ion

O22–

peroxide ion

2O2–super

oxide ion

• Lithium

when it is burnt in air, it forms normal oxide only



C180

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

• Sodium

when it is burnt in an abundant supply of oxygen

forms both the normal oxide and the peroxide



C180

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

e C300

4Na2O(s) + O2(g) 2Na2O2

(s) sodium peroxi

de

• Potassium, rubidium and caesium

form the normal oxide, the peroxide and superoxides when they are b

urnt in air



• Potassium:

4K(s) + O2(g) 2K2O(s)potassium oxide

2K2O(s) + O2(g) 2K2O2(s)potassium peroxide

K2O2(s) + O2(g) 2KO2(s)potassium superoxide



• Rubidium:

4Rb(s) + O2(g) 2Rb2O(s)

2Rb2O(s) + O2(g) 2Rb2O2(s)

Rb2O2(s) + O2(g) 2RbO2(s)



• Caesium:

4Cs(s) + O2(g) 2Cs2O(s)

2Cs2O(s) + O2(g) 2Cs2O2(s)

Cs2O2(s) + O2(g) 2CsO2(s)




Oxides formed by Group I elementsGroup I element

Normal oxide Peroxide Superoxide

Li

Na

K

Rb

Cs

Li2O

Na2O

K2O

Rb2O

Cs2O

–

Na2O2

K2O2

Rb2O2

Cs2O2

–

–

KO2

RbO2

CsO2

• Lithium

does not form the peroxide or superoxide

the size of lithium ion is very small

leading to its high polarizing power



• When a peroxide ion or superoxide ion approaches a lithium ion

the electron cloud of the peroxide ion or superoxide ion would be greatl

y distorted by the lithium ion





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

• The greater the distortion of the electron cloud

the lower the stability of the compound

lithium peroxide and lithium superoxide do not exist



• Potassium ion, rubidium ion and caesium ion

larger sizes

relatively low polarizing power



• The electron cloud of the peroxide ion or superoxide ion

not be seriously distorted by the metallic cations

pack around them with a higher stability

able to form stable peroxides or superoxides



• Beryllium, magnesium and calcium

form normal oxides only on burning in air

2Be(s) + O2(g) 2BeO(s)

2Mg(s) + O2(g) 2MgO(s)

2Ca(s) + O2(g) 2CaO(s)


2. Group II Elements2. Group II Elements

• Strontium and barium

able to form normal oxides and peroxides of the formula MO2 whe

n the metals are burnt in air



• Strontium

2Sr(s) + O2(g) 2SrO(s)strontium oxide



2SrO(s) + O2(g) 2SrO2(s) strontium peroxide

• Barium

2Ba(s) + O2(g) 2BaO(s) barium oxide



2BaO(s) + O2(g) 2BaO2(s)

barium peroxide

500C

700C

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




Oxides formed by Group II elements

Group II element

Normal oxide Peroxide Superoxide

Be

Mg

Ca

Sr

Ba

BeO

MgO

CaO

SrO

BaO

–

–

–

SrO2

BaO2

–

–

–

–

–

• Beryllium peroxide does not exist

the small size of beryllium ion

high polarizing power of beryllium ion



• The small beryllium ion

greatly distorts the electron cloud of the peroxide ion

results in instability of the compound



• Beryllium ion

smaller in size

higher charge than lithium ion



• Beryllium ion

distorts the electron cloud of the peroxide ion to a greater extent th

an lithium ion

beryllium peroxide is very unstable



• The ionic radii of potassium ion and barium ion

very similar



• Potassium

forms stable superoxide on burning in air



• Barium

does not forms stable superoxide



• Barium ion

higher charge than potassium ion

higher polarizing power than potassium ion



• The electron cloud of the superoxide ion

greatly distorted by barium ion



• Barium ion

difficult to pack with the large superoxide ions in a stable ion

ic lattice




Formation of HydroxidesFormation of Hydroxides

• All Group I elements (except lithium)

react with water to form metal hydroxides and hydrogen gas


2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

2K(s) + 2H2O(l) 2KOH(aq) + H2(g)

2Rb(s) + 2H2O(l) 2RbOH(aq) + H2(g)

2Cs(s) + 2H2O(l) 2CsOH(aq) + H2(g)



• They are basic oxides

react exothermically with water to form the corresponding hydrox

ides



• For normal oxides

M2O(s) + H2O(l) 2MOH(aq)



• For peroxides

M2O2(s) + 2H2O(l) 2MOH(aq) + H2O2(a

q)



• For superoxides

2MO2(s) + 2H2O(l) 2MOH(aq) + H2O2(aq) + O2(g)

where M2O, M2O2 and MO2 represent the normal oxides, peroxides and superoxides of Group I elements respectively



• Hydroxides of the Group I elements

the strongest bases known (except lithium hydroxide)



• All Group II elements (except beryllium and magnesium)

react with water to form metal hydroxides and hydrogen gas

less vigorous than the Group I elements in the same period



Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)

Sr(s) + 2H2O(l) Sr(OH)2(aq) + H2(g)

Ba(s) + 2H2O(l) Ba(OH)2(aq) + H2(g)



• Beryllium

does not react with water or steam



• Magnesium

reacts very slowly with water

but reacts more quickly with steam to form magnesium oxide and hydrogen gas

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

(g)



• Calcium and strontium

react readily with water at room temperature and pressure

the reactivity of Group II elements with water increases down the gro

up



• The hydroxides of calcium, strontium and barium

prepared by reacting the normal oxides with water

CaO(s) + H2O(l) Ca(OH)2(aq)

SrO(s) + H2O(l) Sr(OH)2(aq)

BaO(s) + H2O(l) Ba(OH)2(aq)



• Magnesium oxide

only slightly soluble in water

but it dissolves in acids to form salts



• Beryllium oxide

almost insoluble in water or in acids




Calcium reacts readily with water at room temperature and pressure


Ionic Bonding with FixedIonic Bonding with FixedOxidation State in their CompoundsOxidation State in their Compounds

• s-Block elements form compounds

predominantly ionic in nature

show constant oxidation states of +1 for Group I elements and +2 for

Group II elements

• For Group I elements

form ions with an oxidation state of +1 only

their atoms have only one outermost shell electron



• Once this outermost shell electron is removed

a stable fully-filled electronic configuration is obtained

the first ionization enthalpies of Group I elements are low



• The second ionization

involves the removal of an electron from an inner electron shell



• Once this electron is removed

the stable electronic configuration will be disrupted

their second ionization enthalpies are very high



• Group I elements

form predominantly ionic compounds with non-metals

by losing their single outermost shell electrons

form ions having a fixed oxidation state of +1




Chemical formulae of some Group I compounds and the oxidation states of Group I elements in the

compoundsGroup I element

Oxide Hydride ChlorideOxidation state of Group I element in the compound

Li

Na

K

Rb

Cs

Li2O

Na2O2

KO2

RbO2

CsO2

LiH

NaH

KH

RbH

CsH

LiCl

NaCl

KCl

RbCl

CsCl

+1

+1

+1

+1

+1

• For Group II elements

they form ions with an oxidation state of +2 only

their atoms have two outermost shell electrons in the s orbital



• Once these two outermost shell electrons are removed

a stable fully-filled electronic configuration is obtained

the sum of the first and second ionization enthalpies of Group II

elements is relatively low



• The third ionization

corresponds to the removal of an electron from an inner fully-filled

electron shell



• The third ionization enthalpy

an extremely large positive value for these elements

these elements do not form ions with an oxidation state of +3




form predominantly ionic compounds with non-metals

by losing their two outermost shell electrons

form ions having a fixed oxidation state of +2




Chemical formulae of some Group II compounds and the oxidation states of Group II elements in the

compoundsGroup II element

Oxide Hydride ChlorideOxidation state of Group II element in the compound

Be

Mg

Ca

Sr

Ba

BeO

MgO

CaO

SrO

BaO

BeH2

MgH2

CaH2

SrH2

BaH2

BeCl2

MgCl2

CaCl2

SrCl2

BaCl2

+2

+2

+2

+2

+2


Characteristic Flame Colours of SaltsCharacteristic Flame Colours of Salts

• Most s-block elements

give a characteristic flame colour in the flame test



• Method:

putting a sample of the elements or their compounds into a non-lu

minous Bunsen flame



• The outermost shell electrons of atoms of both Groups I and II elements

weakly held by the nucleus

the electrons are easily excited to higher energy levels upon heating



• When these electrons return to their ground states

radiation is emitted

falls into the visible light region of the electromagnetic spectrum



• The amount of energy of the emitted radiation is quantized

the flame colour is a characteristic property of the element



• Example:

Sodium chloride



• When sodium chloride is heated in a Bunsen flame

the ions are converted to gaseous atoms

Na+ Cl–(g) Na(g) + Cl(g)



• The electrons in the gaseous sodium atoms

excited to higher energy levels



• When the excited electrons return to their ground state

light of golden yellow colour is given out

gives a golden yellow flame on burning


The characteristic flame colours of some Groups I and II elements in the flame test

Group I element Flame colour Group II element Flame colour

Li

Na

K

Rb

Cs

Deep reed

Golden yellow

Lilac

Bluish red

Blue

Ca

Sr

Ba

Brick-red

Blood-red or

crimson

Green


Characteristic flame colours of some Groups I and II elements in the flame test: (a) lithium-containing compounds give a deep red flame; (b) sodium-containing compounds give a golden yellow flame; (c) potassium-containing compounds give a lilac flame; (d) calcium containing compounds give a brick-red flame

(a)

(b)

(c)

(d)

Weak Tendency to Form ComplexesWeak Tendency to Form Complexes


A complex is formed when a central metal atom or ion is surrounded by other molecules or ions (called ligands) which form dative covalent bonds with the central metal atom or ion.


• Complex formation

a common characteristic of d-block metal ions



• When a d-block metal ion is surrounded by ligands (such as NH3 and Cl–)

the lone pair electrons of the ligands can be donated to the central d-bl

ock metal ion

form dative covalent bonds



• Example:



• The presence of low-lying vacant d-orbitals in the d-block metal ions

accept the lone pair electrons from the surrounding ligands



• s-Block metal ions

also be surrounded by polar molecules

but there is only electrostatic attraction between the metal io

n and the negative ends of the dipoles



• s-Block metal ions

do not have low-lying vacant orbitals available for forming dative covale

nt bonds

rarely form complexes


Check Point 40-1Check Point 40-1

Variation in Physical PropertiesVariation in Physical Properties

40.2 Variation in Properties of the s-Block Elements (SB p.49)

1. Atomic Radius and Ionic Radius1. Atomic Radius and Ionic RadiusThe atomic radii and ionic radii of most Groups I

elementsGroup I element Atomic radius (nm) Ionic radius (nm)

Li

Na

K

Rb

Cs

0.152

0.186

0.231

0.244

0.262

0.060

0.095

0.133

0.148

0.169


The atomic radii and ionic radii of most Groups II elements

Group II element Atomic radius (nm) Ionic radius (nm)

Be

Mg

Ca

Sr

Ba

0.112

0.160

0.197

0.215

0.217

0.031

0.065

0.099

0.113

0.135


Variations in atomic radius and ionic radius of Groups I and II elements


• Atoms of all Groups I and II elements

form their respective M+ and M2+ ions by losing their outermost s electro

ns

1.1. The ionic radius of any Groups I or II The ionic radius of any Groups I or II element is smaller than its atomic raelement is smaller than its atomic radiusdius

• There is one electron shell less in the cation than the atom

the nucleus pulls the electron cloud more closely towards it

the ionic radius is smaller than the atomic radius

1.1. The ionic radius of any Groups I or II The ionic radius of any Groups I or II element is smaller than its atomic raelement is smaller than its atomic radiusdius


• The atoms or the ions

more electron shells occupied

the outermost electron shells become further away from the nucleus

2.2. Going down both Groups I and II, boGoing down both Groups I and II, both the atomic radii and ionic radii inth the atomic radii and ionic radii increasecrease


• The outermost shell electrons

more effectively shielded by the inner shell electrons from the nuclear

charge



• A decrease in the attractive force between the nucleus and the outermost shell electrons

both the atomic radii and ionic radii increase down a group



• Each atom of the Group II element

one more electron in the outermost shell

one more proton in the nucleus than each atom of the Group I element

in the same period

3.3. Going from Group I to Group II in eaGoing from Group I to Group II in each period, the atomic radii and ionic ch period, the atomic radii and ionic radii decreaseradii decrease


• The additional electron

enters the same shell

at approximately the same distance from the nucleus



• The repulsion between electrons

relatively ineffective to cause an increase in both the atomic radius

and ionic radius



• There is only very little or even no increase in the shielding (or screening) effect of the inner shell electrons on the outermost shell electrons



• The additional proton in the nucleus of the atoms of Group II elements

stronger attractive force to the electrons

the atomic radii and ionic radii of the Group II elements are smaller



2. Ionization Enthalpy2. Ionization Enthalpy


Ionization enthalpies of Groups I elements

Group I elementFirst ionization

enthalpy (kJ mol–1)Second ionization enthalpy (kJ mol–1)

Li

Na

K

Rb

Cs

Fr

519

494

418

402

376

381

7 300

4 560

3 070

2 370

2 420

–


Ionization enthalpies of Groups II elements

Group II element

First ionization enthalpy (kJ mol–

1)

Second ionization enthalpy (kJ mol–

1)

Third ionization enthalpy (kJ mol–1)

Be

Mg

Ca

Sr

Ba

Ra

900

736

590

548

502

510

1 760

1 450

1 150

1 060

966

979

14 800

7 740

4 940

4 120

3 390

–


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


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


• The outermost s electron

located in a new electron shell

1.1. The first ionization enthalpies of GroThe first ionization enthalpies of Group I elements are relatively lowup I elements are relatively low


• The attractive force between thiss electron and the nucleus

relatively weak




effectively shielded from the attraction of the nucleus by the full

y- filled inner electron shells



• Once this electron is removed

a stable octet or duplet electronic configuration is attained




relatively easy to be removed

the first ionization enthalpies for Group I elements are relatively

low



• The second ionization of Group I elements

involves the loss of an inner shell electron

closer to the nucleus



• The removal of this electron

disrupts the stable electronic configuration

the second ionization enthalpies of Group I elements are extremely high



• The two outermost s electrons of atoms of Group II elements

relatively easy to be removed

these two electrons are effectively shielded from the nucleus by i

nner electron shells

2.2. The first and second ionization enthThe first and second ionization enthalpies of Group II elements are relatialpies of Group II elements are relatively lowvely low


• Once these two electrons are removed

a stable octet or duplet electronic configuration is attained



• The third ionization of Group II elements

involves the loss of an inner shell electron

closer to the nucleus



• The stable electronic configuration

disrupte after the removal of this electron

the third ionization enthalpies of Group II elements are very hig

h



• There is an increase in atomic radius down the groups

3.3. Going down both Groups I and II, thGoing down both Groups I and II, the ionization enthalpy generally decre ionization enthalpy generally decreaseseases


• The outermost shell electrons of the atoms

far away from the nucleus

experience less attraction from the positively charged nucleus

less energy is required to remove the outermost shell electrons of the atoms

3.3. Going down both Groups I and II, thGoing down both Groups I and II, the ionization enthalpy generally decre ionization enthalpy generally decreaseseases


3. Hydration enthalpy 3. 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 = H

hyd

always has a negative value



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

hyd

the amount of energy released resulting from the attraction betwe

en ions and water molecules

depends on the charge to radius ratio of the ion



• The greater the charge on the ion

the stronger the attraction between the ion and the water molecule

s

the larger the amount of energy given out



• The smaller the size of the ion

the higher the effective nuclear charge

the stronger the attraction between the ion and the water molecule

s

a large amount of energy is released



Hydration enthalpies of the ions of Groups I and II elements

Group I ionHydration enthalpy

(kJ mol–1)Group II ion

Hydration enthalpy (kJ mol–1)

Li+

Na+

K+

Rb+

Cs+

Fr+

–519

–406

–322

–301

–276

–

Be 2+

Mg2+

Ca2+

Sr2+

Ba2+

Ra2+

–2 450

–1 920

–1 650

–1 480

–1 360

–


Variations in hydration enthalpy of the ions ofGroups I and II elements


1.1. Going down both Groups I and II, thGoing down both Groups I and II, the hydration enthalpy of the ions decre hydration enthalpy of the ions decreases (becomes less negative)eases (becomes less negative)

• The ions

larger in size down both groups

the charge density of the ions falls



• The electrostatic attraction between the ions and water molecules

weaker



• The hydration enthalpy

less negative on going down the groups


2.2. The ions of Group II elements have The ions of Group II elements have more negative hydration enthalpies tmore negative hydration enthalpies than those of Group I elements in the han those of Group I elements in the same periodsame period

• Group II ions

a charge of +2

a smaller ionic radius

higher charge density than the Group I ions in the same period


2.2. The ions of Group II elements have The ions of Group II elements have more negative hydration enthalpies tmore negative hydration enthalpies than those of Group I elements in the han those of Group I elements in the same periodsame period

• The electrostatic attraction between the Group II ions and water molecules

stronger than that between the Group I ions and water molecules


• The melting points of Groups I and II elements depend on

the strength of the metallic bonding

how the atoms are arranged in the metallic crystal lattice

4. Melting Point4. Melting Point


• The stronger the metallic bond

the higher the melting point of the element



• The strength of metallic bond depends on:

1. the ionic radius of the metal ion;

2. the number of valence electrons per atom of the element



The melting points of Groups I and II elements

Group Ielement

Melting Point (C)

Group IIelement

Melting Point (C)

Li

Na

K

Rb

Cs

Fr

180

97.8

63.7

38.9

28.7

24

Be

Mg

Ca

Sr

Ba

Ra

1280

650

850

768

714

697


Variations in melting point of Groups I and II elements


• In each group

the number of electrons participating in metallic bonding remains the same

the ionic radii of the metal ions increase

1.1. Going down both Groups I and II, thGoing down both Groups I and II, the melting points generally decreasee melting points generally decrease


• The attractive forces between the electron sea and the metal ions

weaker

the metallic bond becomes weaker going down the groups

1.1. Going down both Groups I and II, thGoing down both Groups I and II, the melting points generally decreasee melting points generally decrease


2.2. The melting points of Group II elemeThe melting points of Group II elements are much higher than those of Grnts are much higher than those of Group I elementsoup I elements

• The ions of Group I elements

carry one positive charge only

• The ions of Group II elements

carry two positive charges



• The Group II elements

the greater number of valence electrons

the smaller ionic radii



• The attractive forces between the metal ions and the electron sea

stronger in Group II elements



• The metallic bond

stronger in Group II elements


3.3. The general decrease in melting poiThe general decrease in melting point down Group II elements is broken nt down Group II elements is broken with an irregularity with an irregularity ―― the melting p the melting point of magnesium is lower than that oint of magnesium is lower than that of calciumof calcium

• Melting point of a metal depends on

how the individual atoms are packed

in the crystal lattice


3.3. The general decrease in melting poiThe general decrease in melting point down Group II elements is broken nt down Group II elements is broken with an irregularity with an irregularity ―― the melting p the melting point of magnesium is lower than that oint of magnesium is lower than that of calciumof calcium

• Magnesium atoms

not particularly well packed in the

crystal lattice as compared to calci

um atomsCheck Point 40-2ACheck Point 40-2A


Variation in Chemical PropertiesVariation in Chemical Properties

• s-Block elements have strong reducing power

reflected by their low ionization enthalpies



• The lower the ionization enthalpy

the stronger the reducing power of the metal



• Going down both Groups I and II

the atomic size increases

easier for the atoms to lose the outermost shell electrons

the ionization enthalpy decreases accordingly

the reducing power of the metals increases down the groups



• Group I metals

react readily with oxygen and water by losing their single outermos

t shell electron



• Group II metals

generally less reactive than Group I metals

their higher ionization enthalpies


• Most s-block elements

show a silvery white lustre when they are freshly cut

they tarnish rapidly upon exposure to the atmosphere

they react with oxygen in the air to form an oxide layer

1. Reaction with Oxyg1. Reaction with Oxygenen


Sodium shows a silvery white lustre when freshly cut


• s-block elements (except beryllium and magnesium)

very reactive

stored under paraffin oil or in vacuum-sealed ampoules

prevent contact with oxygen and water vapour in the air




form an oxide layer with oxygen

but they tarnish comparatively slowly




burn in air to form one or more of the three types of oxides



• Three types of oxides

normal oxides

peroxides

superoxides



“Dot-and-cross” diagrams of the three types of oxide ions


• On burning in air, lithium

forms only lithium oxide (Li2O)



• On burning in air, sodium

forms a mixture of sodium monoxide (Na2O) and sodium peroxide (Na2

O2)



• On burning in air, potassium, rubidium and caesium

form the normal oxides, peroxides and superoxides




4K(s) + O2(g) 2K2O(s) potassium oxide

C180

• Example:

C300

2K2O(s) + O2(g) 2K2O2(s)potassium peroxi

deK2O2(s) + O2(g) 2KO2(s)

potassium superoxide

C3000



form the normal oxides when they are burnt in air

2Be(s) + O2(g) 2BeO(s)

2Mg(s) + O2(g) 2MgO(s)




relatively unreactive towards oxygen at room temperature and pressure

burn with a brilliant white flame when ignited



Magnesium burns vigorously in air

Magnesium


Magnesium oxide is formed after burning

Magnesium oxide



form the normal oxides and peroxides on burning in air

2Ba(s) + O2(g) 2BaO(s) barium oxide

1. Reaction with 1. Reaction with

OxygenOxygen

2BaO(s) + O2(g) 2BaO2(s)

barium peroxide

500C

700C



very reactive at room temperature and pressure

stored under paraffin oil to prevent contact with air



Oxides formed by the s-block elements

Type of oxide FormulaMetals that form oxides when exposed to air or burnt in air

Normal oxide O2– All Groups I and II elements

Peroxide O22– Na, K, Rb, Cs, Sr, Ba

Superoxide O2– K, Rb, Cs


• All Group I metals

react with cold water to form metal hydroxides and hydrogen

2. Reaction with Water2. Reaction with Water


• Example:

2Li(s) + 2H2O(l) 2LiOH(aq) + H2(g) (vigorous)

2Na(s) + 2H2O(l) 2NaOH(aq) + H2

(g) (violent)

2K(s) + 2H2O(l) 2KOH(aq) + H2(g) (explosi

ve)



• All Group I metals

reduce cold water to form hydroxide ions and hydrogen


H2O(l) + e– OH–(aq) + H2(g)21


• The reactivity of Group I metals with water

related to the relative ease of the metal atoms to lose the outermost

shell electron



• Going down the group

the atomic size increases

the outermost shell electron becomes easier to be removed

the reactivity of Group I metals towards water increases down the

group



• Lithium

reacts with water vigorously



• Sodium

reacts with water violently

moves on the water surface with a hissing sound



• Potassium

reacts with water explosively

producing a lilac flame



• Rubidium and caesium

react with water explosively



Lithium reacts with water vigorously


Sodium reacts with water violently


Potassium reacts with water explosively


• For Group II metals

less reactive than the Group I metals in the same period



• Beryllium

does not react with water



• Magnesium

reacts with steam rapidly to form magnesium oxide and hydrogen g

as

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

(g) (vigorous)



• Magnesium

reacts with cold water very slowly to form magnesium hydroxide and

hydrogen gas

Mg(s) + 2H2O(l) Mg(OH)2(aq) + H2(g)

(very slow)



• Calcium

reacts with water steadily

Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)

(moderate)




react with water vigorously

Sr(s) + 2H2O(l) Sr(OH)2(aq) + H2(g)

(vigorous)


Check Point 40-2BCheck Point 40-2B

40.3 Variation in Properties of the Compounds of the s-Block Elements (SB p.59)

Reactions of Oxides of Reactions of Oxides of ss-Block El-Block Elementsements

• The oxides of all Group I elements

react with water to form the corresponding hydroxides

it is exothermic


• For the normal oxides of Group I elements

react with water to form metal hydroxides

as the only products

Li2O(s) + H2O(l) 2LiOH(aq)



• For peroxides and superoxides of Group I elements

other products are formed



• For peroxides of Group I elements

react with water to form metal hydroxides and hydrogen pero

xide

Na2O2(s) + 2H2O(l) 2NaOH(aq) + H2O2(a

q)




Dissolution of sodium peroxide in water containing phenolphthalein (The red colour is due to the formation of hydroxide ions which turn phenolph

thalein from colourless to red)

• For the superoxides of Group I elements

react with water to form metal hydroxides, hydrogen peroxid

e and oxygen

2KO2(s) + 2H2O(l) 2KOH(aq) + H2O2(aq) + O

2(g)



• The basicity of the oxides of Group I elements increases down the group



• The oxides of Group II elements

generally less basic than those of Group I elements



• The oxides of all Group II elements (except beryllium oxide and magnesium oxide)

react with water to form weakly alkaline solutions



• Example:

CaO(s) + H2O(l) Ca(OH)2(aq) (weakly alkaline)

SrO(s) + H2O(l) Sr(OH)2(aq) (weakly alkaline)



• The basicity of the oxides of Group II elements also increases down the group



• Beryllium oxide

amphoteric

almost insoluble in water or in acids

its amphoteric nature is shown only in its reactions with hot acids or alkal

is



BeO(s) + 2H+(aq) hot

Be2+(aq) + H2O(l)

BeO(s) + 2OH–(aq) + H2O(l) hot

[Be(OH)4]2–(a

q)beryllate io

n



• Magnesium oxide

only slightly soluble in water

dissolves in acids to form salts



• The oxides of other Group II elements

soluble in water

solubility increases down the group



• Barium peroxide

reacts readily with cold water to form barium hydroxide and hydrogen

peroxide

BaO2(s) + 2H2O(l) Ba(OH)2(aq) + H2O2

(aq)



• Metal peroxides and metal superoxides

strong oxidizing agents

indicated by their reactions with water to give hydrogen peroxide



• Sodium peroxide

oxidizes green chromium(III) hydroxide to yellow sodium

chromate(VI)



• Metal peroxides and metal superoxides

useful in both qualitative and quantitative analysis

2Cr(OH)3(s) + 3Na2O2(s) 2Na2CrO4(aq)

+ 2NaOH(aq) + 2H2

O(l)



• All oxides of Groups I and II elements (Except beryllium oxide is amphoteric)

basic

react readily with dilute acids

2. Reaction with Dilute Acids2. Reaction with Dilute Acids


• Their normal oxides

neutralize the acids to form salts and water

CaO(s) + 2HCl(aq) CaCl2(aq) + H2

O(l)



• Their peroxides

react with dilute acids to form salts and hydrogen peroxide

Na2O2(s) + 2HCl(aq) 2NaCl(aq) + H2O2(a

q)



• Their superoxides

react with dilute acids to form salts, hydrogen peroxide and oxyge

n

2KO2(s) + 2HCl(aq) 2KCl(aq) + H2O2(aq) + O2

(g)



• In general, the oxides of s-block elements do not react with dilute alkalis

3. Reaction with Dilute Alkalis3. Reaction with Dilute Alkalis


• Except beryllium oxide

amphoteric

reacts with sodium hydroxide toform sodium beryllate

BeO(s) + 2NaOH(aq) + H2O(l) Na2Be(O

H)4(aq)


3. Reaction with Dilute Alkalis3. Reaction with Dilute Alkalis


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements

Thermal stability refers to the resistance of a compound to undergo decompositionon heating.


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• The higher the thermal stability of a comp

ound

the higher the temperatureneeded for the compound to

decompose thermally


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• A compound is often said to be

thermally stable

not decomposed at the temperature of the normal Bunsen flame

(approximately 1300 K)


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• The thermal stability of an ionic compoun

d depends on

the charges of its constituent ions

the sizes of its constituent ions


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• For compounds with relatively large polar

izable anions

the thermal stability is affected by the polarizing power of the cations


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• If the cation has strong polarizing power

distort the electron clouds of the neighbouring anions to a greater

extent


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• The compound

less stable

more likely to decompose on heating


The large anion is polarized by the small cation


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• Most carbonates and hydroxides of Grou

ps I and II metals

undergo decomposition on heating to give oxides


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• Example:

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

Ca(OH)2(s) CaO(s) + H2O(g)


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• Oxide ions

smaller in size than carbonate ions and hydroxide ions

less polarizable


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• The electrostatic attraction between the c

ations and oxide ions

stronger


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• The oxides of Groups I and II metals

more stable than the corresponding carbonates and hydroxides

decompose to give oxides upon heating


When a compound with large anions undergoes thermal decomposition, a compound with small anions will be

formed since small anions are less easily polarized


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• Group II metal ions

smaller in size

higher charge

stronger polarizing power


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• The large carbonate ions and hydroxide i

ons

distort to a greater extent

more readily to undergo thermal decomposition


Effect of sizes of the cations on thermal stability of the carbonates and hydroxides of both

Groups I and II metals


Relative Thermal Stability of the CarRelative Thermal Stability of the Carbonates and Hydroxides ofbonates and Hydroxides ofss-Block Elements-Block Elements• Going down each group

the size of the cations increases

the polarizing power of the cations decreases

the thermal stability increases


• The carbonates of all Group I metals (except lithium carbonate)

withstand up to a temperature around 800C

1. The Carbonates1. The Carbonates


• Lithium carbonate

decomposes at around 700C

form lithium oxide and carbon dioxide


Li2CO3(s) Li2O(s) + CO2

(g)

C700


• Lithium carbonate

relatively unstable

lithium ion has the smallest size

the polarizing power is the highest



• The electron cloud of any large anion

distort to a great extent

decompose more readily on heating



• Group II metal ions

higher charges

smaller sizes

higher polarizing power



• The electron cloud of the carbonate ion is much distorted

more readily to undergo thermal decomposition



• Example:


C100BeCO3(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


• The hydroxides of all Group I metals (Except lithium hydroxide)

stable to heating with a Bunsen flame

2. The Hydroxides2. The Hydroxides


• Lithium hydroxide

the least stable on heating



• Lithium ion

extremely small

high polarizing power

distorts the electron cloud of the hydroxide ion

decomposition occurs


2LiOH(s) Li2O(s) + H2O(g)


• The hydroxides of Group II metals

less stable to heat

greater the polarizing power




cationic sizes increases

thermal stability increases



• The enthalpy changes

provide evidence for the trend of increasing thermal stability



• Example:


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

ol–1 Mg(OH)2(s) MgO(s) + H2O(g)

H = +81 kJ mol–1


• Example:


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

ol–1 Sr(OH)2(s) SrO(s) + H2O(g)

H = +127 kJ mol–1

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

ol–1


Relative Solubility of the SulphatesRelative Solubility of the Sulphates(VI) and Hydroxides of(VI) and Hydroxides ofs-Block Elementss-Block Elements

1. Processes involved in Dissolution a1. Processes involved in Dissolution and their Energeticsnd their Energetics

• When an ionic solid dissolves in water

two processes are taking place


• Two processes are

1. the breakdown of the ionic lattice

2. the subsequent stabilization of the ions by water molecules (this

process is called hydration)



• When an ionic solid dissolves in water

there must be energetically favourable interactions between t

he water molecules and the dissolved ions



• These interactons

compensate for the breaking of ionic bonds present in the ionic lattice

considered from the point of view of energetics



• The first process

involves an absorption of energy

break down the ionic lattice



• The second process

involves a release of energy

the ions are hydrated



• Example:

The dissolution of sodium chloride in water



• The enthalpy change of solution of sodium chloride = +4 kJ mol–1

when one mole of sodium chloride is completely dissolved in a sufficien

tly large volume of solvent to form an infinitely dilute solution



• NaCl(s) Na+(g) + Cl–(g)

1.1. The sodium chloride solid lattice is bThe sodium chloride solid lattice is broken down to give its constituent ioroken down to give its constituent ions in the gaseous statens in the gaseous state


• The enthalpy change

accompanies this process is the reverse of the lattice enthalpy of sodium



• The lattice enthalpy of sodium chloride= –776 kJ mol–1



• The enthalpy change of the above process = +776 kJ mol–1

NaCl(s) Na+(g) + Cl–(g) H = +776 kJ m

ol–1



2.2. The hydration of the resulting ionsThe hydration of the resulting ions

The enthalpy change of hydration is the enthalpy change accompanies the hydration of one mole of both of these gaseous ions.

Na+(g) + Cl–(g) Na+(aq) + Cl–(aq) H = –772 kJ mol–1


• According to Hess’s law



• The enthalpy change of solution:

Hsoln = Hhyd – Hlattice



• For a salt to be soluble in water

its enthalpy change of solution has to be a negative or a small positive

value



• The sulphates(VI) and hydroxides of Group I metals

more soluble in water than those of Group II metals

2. Relative Solubility of the Sulphates2. Relative Solubility of the Sulphates(VI) and Hydroxides(VI) and Hydroxides


• Group I metal ions

larger sizes

smaller charge



• The Hlattice of their compounds

smaller in magnitude than those of Group II compounds

the dissolution of Group I compounds is more exothermic

the enthalpy changes of solution are more negative



• For the sulphates(VI) of Group II metals

the cations are much smaller than theanions



• The Hlattice is mainly determined by


the reciprocal of the sum of cationic

and anionic radii (i.e. ) rr1


• The large ionic radius of the anion

the much smaller sizes of cations

relatively insignifiant in contributing to the sum of r + and r –




the increase in size of the cations does not cause a significant chan

ge in the Hlattice



• The increase in size of the cations

cause the Hhyd to become less negative



• That is to say

the decrease in Hhyd is more significant than the decrease in Hlattice

the Hsoln becomes less negative

the solubility of the sulphates(VI) of Group II metals decreases



• For the hydroxides of Group II metals

the sizes of anions and cations are of the same order of magnitude



• Again


rr1

the Hlattice is proportional to



cationic size increases

the change in Hhyd is comparatively small

less energy is required to break down the ionic lattice (i.e. the Hlattice

becomes less negative)



• That is to say

the decrease in Hlattice is more significant than the decrease in

Hhyd



• The Hsoln

becomes more negative

the solubility of the hydroxides of Group II metals increases


Check Point 40-3Check Point 40-3

The END

Metals are sometimes referred to as electropositive elements. Why?

AnswerThey have low electronegativity values.

Back


s-Block compounds give a characteristic flame colour in the flame test. Based on this, can you give one use of

s-block compounds?Answer

s-Block compounds can be used in fireworks.

Back


(a) Which ion has a greater ionic radius, potassium ion or calcium ion? Explain your answer.

Answer(a) Potassium ion (0.133 nm) has a greater ionic radius than calcium i

on (0.099 nm) . In fact, potassium ion and calcium ion are isoelect

ronic and have the same number of electron shells. However, calc

ium ion has one more proton than potassium ion, the electron clou

d of calcium ion will experience greater attractive forces from the n

ucleus. This leads to a smaller ionic radius of calcium ion.


(b) Explain why Group I elements show a fixed oxidation state of +1 in their compounds in terms of ionization enthalpies. Answer


(b) Group I elements form ions with an oxidation state of +1 only. It is

because they have only one outermost shell electron. Once this o

utermost shell electron is removed, a stable fully-filled electronic c

onfiguration is obtained. Therefore, the first ionization enthalpies

of Group I elements are low. The second ionization involves the r

emoval of an electron from an inner electron shell. Once this elect

ron is removed, the stable electronic configuration will be disrupte

d. Therefore, their second ionization enthalpies are very high. As

a result, Group I elements form predominantly ionic compounds

with non-metals by losing their single outermost shell electron, an

d they form ions having a fixed oxidation state of +1.


(c) Ions of Group I and Group II elements have a very low tendency to form complexes. Give one reason to explain your answer. Answer

(c) As ions of Group I and Group II elements do not have low-lying vaca

nt orbitals available for forming dative covalent bonds with the lone p

air electrons of surrounding ligands, they rarely form complexes.


(d) Give one test which would enable you to distinguish a sodium compound from a potassium compound.

Answer(d) Sodium compounds and potassium compounds can be distinguishe

d by conducting a flame test. In the flame test, sodium compounds

give a golden yellow flame, while potassium compounds give a lilac

flame.

Back


What is a dative covalent bond? How is it formed?

AnswerA dative covalent bond is a covalent bond in which the shared pair o

f electrons is supplied by only one of the bonded atoms. A dative co

valent bond is formed by the overlapping of an empty orbital of an at

om with an orbital occupied by a lone pair of electrons of another ato

m.

Back


(a) (i) List the factors that affect the value of the ionization enthalpy of an atom.

Answer(a) (i) There are four main factors affecting the magnitude of the

ionization enthalpy of an atom. They are the electronic configu

ration of an atom, the nuclear charge, the screening effect, an

d the atomic radius.


(a) (ii) Why is ionization enthalpy of an atom always positive?

Answer(a) (ii) Ionization enthalpy of an atom always has a positive value

because energy is required to overcome the attractive forces b

etween the nucleus and the electron to be removed.


(a) (iii) Describe the general trend of the first and second ionization enthalpies down Group I of the Periodic Table. Answer



(a) (iii) The first ionization enthalpies of Group I elements are rela

tively low. The outermost s electron is located in a new electro

n shell. The attractive force between this s electron and the nu

cleus is relatively weak. Also, this s electron is effectively shiel

ded from the attraction of the nucleus by the fully-filled inner el

ectron shells. Once this electron is removed, a stable octet or

duplet electronic configuration is obtained. Consequently, this

s electron is relatively easy to be removed, and hence the first

ionization enthalpies of Group I elements are relatively low. Ho

wever, the second ionization of Group I elements involves the l

oss of an inner shell electron which is closer to the nucleus. Th

e removal of this electron disrupts the stable electronic configu

ration. Therefore, the second ionization enthalpies of Group I e

lements are extremely high.

(b) (i) List the factors that affect the value of the hydration enthalpy of an ion.

Answer(b) (i) The value of the hydration enthalpy of an ion depends on

the size and the charge of the ion.


(b) (ii) Why does hydration enthalpy of an ion always have a negative value?

Answer


(b) (ii) Hydration enthalpy of an ion always has a negative value

because it is the amount of energy released resulting from the

attraction between the ion and water molecules.

(b) (iii) Describe the general trend of the hydration enthalpy down Group II of the Periodic Table.

Answer


(b) (iii) Going down Group II, the hydration enthalpy of the ions d

ecreases (becomes less negative). Since the ions get larger i

n size on moving down the group, the charge density of the io

ns falls. As a result, the electrostatic attraction between the io

ns and water molecules becomes weaker, and the hydration

enthalpy becomes less negative down the group.

Back

The burning of lithium, sodium and potassium in oxygen gives different types of oxides. Why do the metals beh

ave differently?Answer


On burning in air, lithium forms only lithium oxide, and it does not form the

peroxide or superoxide. This is because the size of lithium ion is very small,

leading to its high polarizing power. When a peroxide ion or superoxide ion

approaches a lithium ion, the electron cloud of the peroxide ion or superoxi

de ion (large in size) would be greatly distorted by the lithium ion. The great

er the distortion of the electron cloud, the lower the stability of the compoun

d. That is why lithium peroxide and lithium superoxide do not exist. Sodium

ion has a larger size than lithium ion. Its lower polarizing power allows it to f

orm the peroxide when sodium is burnt in air. Potassium ion has a much lar

ger size, so it has relatively low polarizing power. The electron cloud of the

peroxide ion or superoxide ion would not be seriously distorted by potassiu

m ion. This allows the peroxide ions or superoxide ions to pack around pota

ssium ion with a higher stability. As a result, potassium is able to form stabl

e peroxide or superoxide on burning in air.

Back


(a) Suggest a reason why the reaction of lithium with water is less vigorous than those of sodium and potassium.

Answer(a) The reactivity of Group I metals with water is related to the relative

ease of the metal atoms to lose the outermost shell electron. Going

down the group, as the atomic size increases, the outermost shell e

lectron becomes easier to be removed. Therefore, the reactivity of

Group I metals towards water increases down the group. Lithium re

acts with water vigorously. Sodium reacts with water violently and

moves on the water surface with a hissing sound.


(b) Which element is the strongest reducing agent, calcium, strontium or barium?

Answer(b) Barium is the strongest reducing agent. It is because the reducing p

ower of an element is related to the ease of the atom to lose the out

ermost shell electron. Since barium has larger atomic sizes, its oute

rmost shell electrons are less firmly held by the nucleus. Therefore,

barium has a higher tendency to lose its outermost shell electrons t

han both calcium and strontium.


Back

The value of Hsoln of a solid does not indicate whether the solid is soluble in water or not. So how can we predic

t the solubility of a solid in water?Answer

Back


Generally speaking, for a solid to be soluble in water, its enthalpy chan

ge of solution has to be a negative or a small positive value.

(a) Give balanced chemical equations for the following reactions:

(i) Thermal decomposition of barium carbonate

(ii) Reaction between sodium peroxide and water

(iii) Reaction between calcium oxide and dilute hydrochloric acid

Answer


(a) (i) BaCO3(s) BaO(s) + CO2(g)

(ii) Na2O2(s) + 2H2O(l) 2NaOH(aq) + H2O2(aq)

(iii) CaO(s) + 2HCl(aq) CaCl2(aq) + H2O(l)

(b) Suggest a reason why barium sulphate(VI) is insoluble in water, while potassium sulphate(VI) is soluble inwater although they have cations of similar sizes and the same anion.

(The ionic radii of potassium ion and barium ion are 0.133 nm and 0.135 nm respectively.) Answer


(b) When an ionic solid dissolves in water, two processes are taking pla

ce. They are the breakdown of the ionic lattice and the subsequent s

tabilization of the ions by water molecules. The enthalpy change inv

olved in the whole dissolution process is known as the enthalpy cha

nge of solution, Hsoln, which is equal to Hsoln = Hhyd – Hlattice. For

an ionic compound to be soluble in water, the enthalpy change of so

lution has to be a negative or a small positive value. The reason why

barium sulphate(VI) is insoluble in water while potassium sulphate(V

I) is soluble in water is that potassium ion has a smaller charge than

barium ion. The Hlattice of potassium sulphate(VI) is smaller in magni

tude (less negative) than that of barium sulphate(VI). As a result, the

enthalpy change of solution of potassium sulphate(VI) is more negat

ive, and hence it is soluble in water while barium sulphate(VI) is not.


(c) Compare the solubility of calcium sulphate(VI) and barium sulphate(VI) in water. Explain your answer.

Answer(c) Calcium sulphate(VI) is expected to be more soluble than barium sul

phate(VI). It is because calcium ion has a smaller size than barium i

on. This causes the Hhyd of calcium sulphate(VI) to be more negativ

e than that of barium sulphate(VI). As a result, the Hsoln of calcium s

ulphate(VI) becomes more negative than that of barium sulphate(VI),

and hence calcium sulphate(VI) is more soluble in water than bariu

m sulphate(VI).


Back

Date post:	13-Jan-2016
Category:	Documents
Upload:	godfrey-allison
View:	339 times
Download:	5 times

The s-Block Elements 40.1Characteristic Properties of the s-Block Elements 40.2Variation in...

Documents