The sBlock Element

1
Section 9

• The s-Block Element

2
Group I Elements (Alkali Metals)

• They have similar chemical properties. They are
  soft metals with fixed O.N. 1 in their
  compounds.

3
Group II Elements (Alkaline Earth Metals)

• Mg and Ca are the most abundant elements and Ra
  is the most scarce element which is unstable and
  radioactive.

4
Atomic and ionic radius

• Atomic radius and ionic radius increase
• An addition of one more shell ENC decrease

5
?1 Compare the atomic radius of alkali and
alkaline earth metals

• Alkali metals gt alkaline earth metals because
• Increase in nuclear charge gt increase in
  shielding effect
• ENC increases and electron experience a larger
  nuclear attraction.

6
?2 Compare the atomic and ionic radius

• Ionic radius lt corresponding atomic radius
• Same nuclear charge, weaker shielding effect
• ENC increases, stronger nuclear attraction
  towards electrons.

7
Melting and boiling point

• M.P. decreasedENC decreased so that nuclear
  attraction towards electrons ( metallic bond
  strength)decreased.

8
?4. Difference between m.p. and b.p. of s-block
metals

• b.p. gt m.p. a lot
• Most of the metallic bonds remains in the liquid
  state nearly all bonds in liquid state have to
  be broken on vaporization.

9
Atomic volume

• Atomic volume increasedAtomic size increases,
• Metallic bond strength decreases.

10
?5. Compare the atomic volumes of s-block metals

• Alkali metals gt Alkaline earth metals
• Atomic size of alkaline earth metals lt alkali
  metals
• Metallic bond strength of alkaline earth metals gt
  alkali metals

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Density

• Density increasedAtomic mass increased to a
  greater extent than atomic volume.

12
?6. Compare the density of s-block metals

• Alkali metals lt Alkaline earth metals
• Atomic mass of alkaline earth metals gt alkali
  metals
• Atomic volume of alkaline earth metals lt alkali
  metals
• Density atomic mass / atomic volume

13
Ionization energy

• I.E. decreasedENC decreases due to addition of
  shells.
• I.E. of alkaline earth metals gt alkali metals
• ENC increases

14
?7. Difference between 1st and 2nd I.P. of alkali
metals

• Electron being removed is from the inner shell
  the electron thus experience a larger nuclear
  attraction. Besides, ENC of an ion would be much
  greater than the corresponding atom as the
  shielding effect is weaker.

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Summary

• On passing down the group
• Atomic radius and ionic radius increasedAn
  addition of one more shell ENC decrease
• I.E. decreasedENC decreased due to addition of
  shells.
• M.P. decreasedENC decreased so that nuclear
  attraction towards electrons decreased.
• Density increasedatomic mass increased to a
  greater extent than atomic volume.
• E.A. decreasedENC decreased so that tendency to
  accept e- decreased.
• Reducing power and reactivity increasedI.E
  decrease and reduction potential become more
  negative.
• Enthalpy of hydration of cation less
  negativeElectrostatic interaction between the
  polar water molecules and ions become less as the
  ionic radius increases.

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Summary of physical properties of Group I and IIA
elements
Atomic and ionic radii atomic volume
m.p. and b.p. density and I.P.
m.p. and b.p. density and I.P.
Atomic and ionic radii atomic volume
17
Variaiton in chemical properties

• Owing to the low value of 1st I.P., alkali metals
  are relatively more easily to form X, and the
  resulting compound is quite stable. The sum of
  1st and 2nd I.P. of alkaline earth metals is not
  too low, yet the lattice energy recorded on
  forming the ionic compounds is large enough for
  the formation of X2.

18
Reducing power and reactivity of s-block elements

• Try to account for the following redox potentials

 
Eo / V
Li(s) ? Li(aq) e
3.04V
Cs(s) ? Cs(aq) e
2.93V
Rb(s) ? Rb(aq) e
2.93V
K(s) ? K(aq) e
2.92V
Na(s) ? Na(aq) e
2.71V
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Reducing power and reactivity of s-block elements

• The redox potential
• M(s) ? M(aq) e depends on
• 1. the formation of separate atom from crystal
  lattice
• M(s) ? M(g)
• ?Hsub heat of sublimation (ve)
• 2. the formation of gaseous ion from gaseous atom
• M(g) ? M(g) e
• I.E. ionization energy (ve)
•  3. the formation of hydrated ion from gaseous
  ion
• M(g) aq ? M(aq)
• ?Hhyd hydration energy (-ve)

20
Reducing power and reactivity of s-block elements

• The overall enthalpy change
• (?H) ?Hsub I.E. ?Hhyd
•  
• ?H more -ve
• ? the greater the redox potential i.e. the
  stronger the reducing agent.

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Reducing power and reactivity of s-block elements

• From Na to Cs, the reduction potential increased.
• But Li has greatest reduction potential.
•  
• On passing down the group, both ?Hatm, I.E.
  decrease but ?Hhyd also become less -ve.
•  
• But ?Hsub and I.E. decrease to a greater extent
  than the ?Hhyd, ?H(overall) is more negative.

22
Reducing power and reactivity of s-block elements

• Li is an exceptional case, it has the greatest
  redox potential. It is because the size of Li
  is very very small (it belongs to 2nd period),
  ?Hhyd is exceptionally more -ve. Therefore
  ?H(overall) is thus more negative.

23
Variation in chemical properties

• Reactions
• With air - All tarnish in air (that is, forming a
  film of oxide on the surface), therefore they are
  stored in paraffin oil.
• When burnt in sufficient amount of oxygen

Kind of oxides
Elements which form this type of oxide in
adequate supply of air
normal oxide O2-
Li, Mg, Ca, Sr
peroxides
Na, Ba
superoxide O2-
K, Rb, Cs
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Reaction with air

• Dot and cross diagram for oxide O2- ion and
  peroxide O22- ion

O O bond can be easily broken
Size of peroxide ion gt size of oxide ion
25
Reaction with air

• Li ion is extremely small, it is not possible
  for sufficient number of peroxide ions to
  surround the Li ion with causing repulsion
  between the anions, therefore only normal oxide
  exists.
• The larger peroxide and superoxide anions are
  stabilized by larger cations due to limiting
  radius ratio.

26
?9. Why Group IIA elements form normal oxides,
except barium ?

• Stability of oxide ion gt peroxide and superoxide
  ion
• Ba2, being the largest ion, has the weakest
  polarizing power electron cloud of the peroxide
  ion will be distorted by other group IIA AND
  also Li metal ions and become unstable.
• Ba2 is the biggest ion in Group IIA slightly
  larger than K, no severe repulsion would occur
  between these large peroxide ion when surrounding
  Ba2 in the lattice.

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Reaction with air

• In case of Li, Mg, Ca, Sr and Ba, the final
  products will be a mixture of nitrides,
  carbonates together with the oxides.
• Only Li in group IA would form Li3N (lithium
  nitride).
• N3- ion is hard to form, why ?
• 6 Li N2 ?2 Li3N
• Li3N 3H2O ? 3LiOH NH3

28
Reaction with water

• With water - All (except Be) reacts to give out
  hydrogen.
• 2H2O 2e- ? 2OH- H2 at pH 7 E
  -0.41V
• 2M 2H2O ? 2MOH 2H2 ?E ve
• (spontaneous)
• But the vigor of reaction K gt Na gt Li although
  ?E of Li is greatest. Why?
• ?E shows the equilibrium position (i.e. the
  reaction is spontaneous or not), ?E increased
  implies that equilibrium lies on the product
  side.
• Rate of reaction must consider the Eact
  (activation energy). From the information given,
  the rate of Li is the slowest among the three.
  That is Eact for Li is the highest, so the rate
  is relatively slow.

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Reaction with water

• The rate of reaction increases on passing down
  the group. Reactivity of alkali metals towards
  water is much higher than alkaline earth metals.
• Magnesium reacts with hot water and steam to give
  magnesium hydroxide and magnesium oxide
  respectively.

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Reaction with acid

• All react vigorously and explosively.
• 2M 2H ? H2 2M
• But reactions between sulphuric acid and
  Ca, Sr, Ba become less vigorous after the
  reaction starts due to the formation of insoluble
  layer of sulphates.
• Ca H2SO4 CaSO4(s) H2

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Reaction with non-metals

• With non-metal - All combine with X2, S and O2, P
  or even H2 at suitable temperature.
• Li, Mg, Ca, Sr and Ba also combine directly with
  nitrogen.
• Ca H2 CaH2 calcium hydride
• 2 K S K2S potassium sulphide
• 3 Mg N2 Mg3N2 magnesium nitride

32
Oxides

• All are white crystalline solid, ionic and
  strongly basic in character. They are hydrolysed
  by water to form corresponding hydroxides. Degree
  of hydrolysis increases down the group, since
  oxide become more ionic.
• O2- H2O 2 OH-

BeO, being exceptional case, is amphoteric BeO
2HCl ? BeCl2 H2O BeO 2OH- H2O
? Be(OH)42- beryllate
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?10. Compare the basic strength of Na2O and MgO

• Na2O is more basic. Na has a weaker
  polarizing power than Mg2 (as the latter one has
  a higher charge/radius ratio OR charge
  density), electron in O2- ion is more available
  to attack hydrogen in water molecule. More
  hydroxide ion is thus formed.

34
Hydrides

• Formed by heating the element in hydrogen gas (at
  400? or above)
• Strong reducing agents
• Hydrolysed by water to form hydrogen gas and
  solution or suspension of hydroxides
• Readiness of hydrolysis increases down the group
  since the hydride is more ionic
• reaction is more vigorous for alkali metals than
  for alkaline earth metals

35
Chlorides

• All are white crystalline solid, soluble in water
  to form hydrated ion.
• NaCl(s) aq. Na(aq) Cl-(aq)
• Hydrated sodium and chloride ion
• No hydrolysis and thus a neutral solution
• But for MgCl2, which is partially ionic,
  hydrolysed by water to give a slightly acidic
  solution.
• MgCl2(s) 6 H2O Mg(H2O)62 2 Cl-(aq)
• Mg(H2O)62 H2O Mg(H2O)5(OH) H3O
• Be(H2O)42 H2O Be(H2O)3(OH) H3O

36
Thermal stability of other compounds

• For a large polarizable anion (e.g. HCO3-,
  CO32-,NO3-, SO42-), the stability depends on the
  polarizing power of the cation. If the cation can
  distort the electron cloud of the anion so much
  that the bonds (e.g. C-O bond in carbonate) is
  weakened, the bond will be easily broken on
  heating to give metallic oxides and gas(es) (CO2
  for carbonate).

37
Thermal stability

• ?
• MgCO3 MgO CO2
• ?
• MgSO4 MgO SO3
• ?
• 2 Mg(NO3)2 2 MgO 2 NO2 O2
• ?
• 2 NaNO3 2 NaNO2 O2

38
?11. Compare the stability of Na2CO3 and MgCO3

• Na2CO3 is thermally stable because polarizing
  power of Na is weaker than Mg2 (as the latter
  one has a higher charge/radius ratio / charge
  density), electron cloud of the carbonate ion is
  much distorted by Mg2 that the C O bond is
  weakened and thus more easily broken when heated

39
Thermal stability

• Polarising power of cation decreases on passing
  down the group as the size of the cation become
  larger. Most group I salts are thermally stable
  except for those of lithium.
• While group II salts are relatively less stable
  to heat. (Note that only lithium carbonate is
  thermally unstable among group I carbonates.).

40
Thermal stability

• Some sodium and potassium salts are decomposed
  when heated
• 2 NaNO3 2 NaNO2 O2
• 2 NaHCO3 Na2CO3 CO2 H2O

41
Solubility of salts in water

• All group I compounds are practically soluble.The
  solubility increases as heat of hydration is more
  negative than lattice energy.

42
Solubility of salts in water

• Lattice energy depends on the sum of the ionic
  radii while the hydration energy depend on ionic
  radius of the individual ions both would
  decrease as size of ions increases. Hydration
  energy of a compound is contributed by both the
  cation and anion.

43
Solubility of salts with large anion (e.g. I-,
SO42-, CO32-) in water

• Hydration energy contributed by anion is small,
  i.e. the hydration energy mainly contributed by
  the cation.
• On passing down the group, hydration energy
  decrease greatly / tremendously.
• The sum of ionic radii only increase slightly as
  the size of anion is large,
• the decrease in lattice energy is small.
• ?Hsoln ?Hhyd ? LE ?
• ?Hsoln become less negative (more positive) on
  passing down the group.

44
Solubility of salts with large anion (e.g. I-,
SO42-, CO32-) in water
45
Solubility of salts with small anion (e.g. F-,
OH-) in water

• Lattice energy decrease more rapidly than the
  hydration energy on passing down the group,
• hydration energy mainly depends on the small
  anion and would not change much
• decrease in lattice energy mainly determines the
  solubility of the salt.
• ?Hsoln ?Hhyd ? LE ?
• ?Hsoln become more negative (less positive) on
  passing down the group.

46
Solubility of salts with small anion (e.g. F-,
OH-) in water
47
Compare the solubility of salts of alkali and
alkaline earth metals

• An increase in charge will increase lattice
  energy to a greater extent than hydration energy,
• salts of alkaline earth metals are generally less
  soluble than that of alkali metals, and
• doubly charged anions give more insoluble
  compounds.

48
General characteristics of s-block elements

• Fixed oxidation state
• The only possible positive oxidation state shown
  by the elements is equal to the total number of
  electron in the outermost shell. This oxidation
  state corresponds to the loss of sufficient
  number of electrons to achieve the octet
  configuration ns2np6, thus only forms compounds
  in which they obtain the octet configuration.
• The loss of more than the valence electron
  requires too much ionization energy, thus
  prevents these metals from showing an oxidation
  number other than the one equal to their group
  number.

49
General characteristics of s-block elements

• Ability to form complexes
• Owing to the lack of underlying (inner) low
  energy vacant orbital, s-block elements rarely
  form complexes.
• The cations which form stable complexes normally
  carrying a high charge / radius ratio, resulting
  in larger electrostatic attraction between the
  central ions and the ligands.

50
Ability to form complexes

• Group I metal ions cannot form hydrated ions of
  definite formula in aqueous solution, though they
  can by hydrated to certain extent.
• Lithium ion, which has the smallest size, show
  certain degree of hydration in crystal of its
  salts.

51
Ability to form complexes

• Group II metals ions has higher charge/radius
  ratio and they have higher tendency to form
  complexes.
• Beryllium forms many complex but barium forms
  very few.
• e.g. BeF3- BeF42- Be(H2O)42
• The most important complex for Mg is chlorophyll,
  which has a very complicated structure with fused
  rings the Mg atom being at the center of the
  rings bonded to 4 nitrogen atoms.

52
Ability to form complexes

• Ca2 and Mg2 form stable complex with strong
  complexing agents. e.g.
• ethylenediaminetetraacetic acid (EDTA) which has
  4 functional oxygen atoms and 2 donor N-atoms per
  molecule.
• More discussion on
• d-block elements.

53
Abnormality of lithium and its compounds among
group IA

• 1 Lithium carbonate and hydroxide are decomposed
  by heat.
• 2 Lithium carbonate, hydroxide and fluoride are
  insoluble in water.
• 3 Lithium forms only normal oxide when reacting
  with
• oxygen.
• 4 Lithium forms nitride when heated in air.
• 5 Lithium ion is highly hydrated in water,
  resulting in lowest
• mobility.
• 6 Lithium hydroxide is not a strong base.
• 7 Almost all lithium salts are hydrated in its
  crystal lattice.

54
Abnormality of lithium and its compounds among
group IA

• Reason
• Exceptional small size of Li ion, e.g. forming
  nitride
• 6 Li(s) N2(g) 2 Li3N(s)
• 6 Li(g) 2 N(g) 2 N3-(g)
• 
  6 Li(g)

The highly negative LE of Li3N offset the energy
required to ionize the nitrogen gas to nitride
ion.
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Diagonal relationship between magnesium and
lithium

• 1. Both only form normal oxide.
• 2. Both give nitrides when heating in air.
• 3. Carbonates,sulphates,hydroxides are decomposed
  by heat to metallic oxides.
• Carbonates, hydroxides are insoluble in water.
• Reasons
• Effective nuclear charge increases on passing
  along the
• period but decreases on passing down the group,
  so Li
• and Mg2 have similar effective nuclear charge
  which in
• turn affecting its polarizing power.

56
Flame Test for Metal Ions