1. S-Block Elements By- Manas Mahajan
ALKALI METALS

2. Chapter summary Introduction
Characteristic properties of the s-block elements .
Variation in properties of the s-block elementsof the First Group(Alkali Metals).
Physical Properties

3. Members of the s-Block ElementsLi Be Na K Rb Cs Fr Mg Ca Sr Ra Ba
IA IIA
IA Alkali metals
IIA Alkaline Earth
metals

4. Characteristic properties of s-block elements
Metallic character
Low electronegativity
Basic oxides, hydroxides
Ionic bond with fixed oxidation states
Characteristic flame colours
Weak tendency to from complex

5. Metallic character High tendency to lose e- to form positive ions
Metallic character increases down both groups

6. Electronegativity Low nuclear attraction for outer electrons
Highly electropositive
Small electronegativity
Group I
Group II Li Be Na
Mg 1.2 K Ca Rb Sr Cs Ba Fr Ra

7. Characteristic flame colours
Na+ Cl- (g)  Na (g) + Cl (g)
Na(g)  Na* (g)
[Ne]3s [Ne]3p1
Na*(g)  Na(g) + h (589nm, yellow)

8. Flame test Li deep red Ca brick red Na yellow Sr blood red K lilac
Rb bluish red
Cs blue
Ca brick red
Sr blood red
Ba apple green
HCl(aq)
sample

9. Variation in properties of elements
Atomic radii
Ionization enthalpies
Hydration enthalpies
Melting points
Reactions with oxygen, water, hydrogen and chlorine

10. Atomic radii (nm) Li 0.152 Be 0.112 Na 0.186 Mg 0.160 K 0.231 Ca 0.197Rb
0.244 Sr
0.215 Cs
0.262 Ba
0.217 Fr
0.270 Ra
0.220 Fr Ra Li Be

11. Ionization Enthapy Group I 1st I.E. 2nd I.E. Li 519 7300 Na 494 4560 K
418
3070 Rb
402
2370 Cs
376
2420
Group I
1st I.E.
2nd I.E.
3rd I.E. Be
900
1760
14800 Mg
736
1450
7740 Ca
590
1150
4940 Sr
548
1060
4120 Ba
502
966
3390

12. Ionization Enthalpy 1st I.E. Be+ Li Na 2nd IE K Rb Ca+ Cs Ba+ Be Ca Ba
300
400
500
600
500
1000
1500
2000 Be Ca Ba
Be+
Ca+
Ba+
1st IE
2nd IE

13. Ionization Enthalpy Group I
Have generally low 1st I.E. as it is well shielded from the nucleus by inner shells.
2. Removal of a 2nd electron is much more difficult because it involves the removal of inner shell electron.
3. I.E. decreases as the group is descended.
As atomic radius increases, the outer e is further away from the well-shielded nucleus.

14. Hydration Enthalpy M+(g) + aqueous  M+(aq) + heat -600 -300 M+
Li+ Na+ K+ Rb+ Cs+ M+

15. Hydration Enthalpy Be2+ Mg2+ Ca2+ Sr2+ Ba2+ -2250 -2000 -1750 -1500
-600
-300
Li+ Na+ K+ Rb+ Cs+

16. Hydration Enthalpy General trends:
On going down both groups, hydration enthalpy
decreases.
(As the ions get larger, the charge density of the
ions decreases, the electrostatic attraction between
ions and water molecules gets smaller.)
Group 2 ions have hydration enthalpies higher
than group 1.
( Group 2 cations are doubly charged and have
smaller sizes)

17. Ionization Energy
Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
First ionization energy is that energy required to remove first electron.
Second ionization energy is that energy required to remove second electron, etc.

18. Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

19. Ionization Enthalpy of Alkali Metals
ionization enthalpy decreases down the group from Li to cs because as we move down a group the number of valence electrons goes increasing separating the electrons away from the nucleus ,there is an increasing shielding of the nuclear charge by the inner shell electrons and thus the removal of electrons requires less energy as we move down.

20. Physical Properties Silvery White Soft & Light Metals
Due to large size , elements have low density
Low Melting & Boling Points indicates weak bonding due to presence of only 1 valence electron .

21. Increasing atomic number
Alkali Metals
In order of increasing atomic number the alkali metals are:
Lithium
Sodium
Potassium
Rubidium
Caesium
Francium
Increasing atomic number

22. Properties of Alkali Metals
All alkali (group 1) metals react violently with water, forming Hydrogen gas and Hydroxides (pH above 7):
Alkali metals are:
Metals found in group 1 of the periodic table.
Soft when cut (compared to other metals).
Metals with low melting points and densities.
Powerful reducing agent and form univalent compounds.
Metals which tarnish in air.

23. Oxides E.g. Li2O Na2O, Na2O2 On cumbustion in excess of air, K2O2, KO2
alkali metals form
Oxides- LiO2
2. Peroxides- Li2O2, NaO2
3. Superoxides- K2O2 ,Cs2O2 RbO2
E.g.
Li2O
Na2O, Na2O2
K2O2, KO2
Rb2O2, RbO2
Cs2O2, CsO2

24. Hydroxides E.g. LiOH NaOH KOH RbOH CsOH
The oxides of the alkali metal are easily
hydrolysed by addition of water.
E.g.:- LiO2 + H2O → LiOH

25. Halides
The alkali metals combine directly with halogens under appropriate conditions forming halides of general formula MX.
These halides can also be prepared by the action of aqueous halogen acids (HX) on metals oxides, hydroxides or carbonate.
All these halides are colourless, high melting crystalline solids having high negative enthalpies of formation.

26. Examples of halides M2O + 2HX → 2MX + H2O MOH + HX → MX + H2O
            
MOH + HX → MX + H2O
       
M2CO3 + 2HX → 2MX + CO2 + H2O
(M = Li, Na, K, Rb or Cs)
(X = F, Cl, Br or I)

27. Salts of oxo-acids Since the alkali metals are highly electropositive,
therefore their hydroxides are very strong bases
and hence they form salts with all oxoacids.
( H2CO3, H3PO4, H2SO4, HNO3, HNO2 etc) .
They are generally soluble in water and stable
towards heat.

28. Li2CO3 , however is considerably less stable and decomposes readily.
   The carbonates (M2CO3)  of alkali metals are
remarkably stable upto 1273 K, above which they first melt and then eventually decompose to form oxides.
Li2CO3 , however is considerably less stable and decomposes readily.
Δ          
Li2CO3 → Li2O + CO2
This is presumably due to large size difference between
Li+ and CO2-3  which makes the crystal lattice unstable.

29. SODIUM CHLORIDE
The most abundant source of sodium chloride is sea water. Crude sodium chloride, generally obtained by crystallisation of brine solution, contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. Calcium chloride,CaCl2, and magnesium chloride MgCl2 are impurities because they are deliquescent(absorb moisture from the atmosphere).

30. To obtain pure sodium chloride–
Crude salt is dissolved in minimum amount of water and filtered to remove insoluble impurities.
The solution is then saturated with hydrogen chloride gas.
Crystals of pure sodium separate out.
Calcium and magnesium chloride, being more soluble than sodium chloride, remains in solution.

31. USES OF NaCl −
It is used as a common salt or table salt for domestic purpose.
It is used for the preparation of Na2O2 , NaOH and Na2CO3.

32. SODIUM HYDROXIDE(CAUSTIC SODA), NaOH
Sodium hydroxide is generally prepared commercially by the electrolysis of sodium chloride in Castner −Kellner cell.
A brine soln. is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam. Chlorine gas is evolved at the anode.

33. NaCl → Na+ + Cl¯ AT ANODE: Cl ¯ ─ e¯ → Cl Cl + Cl → Cl2 AT CATHODE: Na+ + e ¯ → Na+ Na + Hg → NaHg Amalgam The amalgam is treated with water to give sodium hydroxide, mercury and hydrogen gas. 2NaHg + 2H2O→ 2NaOH + 2Hg + H2

34. PROPERTIES- 1. NaOH is a white, translucent solid and it melts at 591K
PROPERTIES- 1. NaOH is a white, translucent solid and it melts at 591K. 2. it is readily soluble in water to give alkaline solution. 3. Crystals of NaOH are deliquescent. It reacts with the CO2 in the atmosphere to form Na2CO3.
USES-
The manufacture of soap, paper and no. of chemicals.
In petroleum refining
In textile industry for mercerising cotton fabrics
As laboratory reagent
For preparation of pure oils and fats

35. SODIUM HYDROGENCARBONATE(BAKING SODA), NaHCO3
Sodium hydrogencarbonate is known as baking soda because it decomposes on heating to generate bubbles of carbon dioxide.
It is made by saturating a solution of sodium carbonate with carbon dioxide. The white crystalline powder of sodium hydrogencarbonate , being less soluble, gets separated out.
Na2CO3 + H2O + CO2 → 2NaHCO3
NaHCO3 is a mild antiseptic for skin infections. It is used in fire extinguishers.

36. BIOLOGICAL IMPORTANCE OF SODIUM AND POTASSIUM
A typical 70 kg man contains about 90g of Na and 170g of K compared with only 5g of iron and 0.06g of copper.
Potassium ions are present in higher concentration inside the cells than sodium ions and they are present outside the cell in blood plasma.
Because of large concentration gradient inside and outside the cells, the transport of sodium ion into the cells is favoured. To pump out these ions again from the cell to maintain concentration gradient large driving force is carried out. The energy for this process is provided by ATP molecules.
Thus both sodium and potassium ions are essential for living organisms.

37. Anomalous properties of Lithium

38. Lithium
Symbol – Li Atomic no. - 3 Atomic Weight – 6.94u Electronic Configuration - 1s22s1 Group no. – 1 Period no. – 2 Group name – Alkali Metals Block name – ‘s’ Standard State(298 K)- Solid Color – Silvery-white/grey Classification - Metallic

39. Anomalous Properties High melting & boiling point.
Much harder than other alkali metals.
Reacts with oxygen least readily to form normal oxide(E.g. Li2O), whereas other alkali metals form peroxides and superoxides(E.g. MO2,M2O2).
4Li + O2 → 2Li2O
Unlike other alkali metals lithium reacts directly with carbon to form an ionic carbide.

40. The carbonates, hydroxides and nitrates of lithium decompose on heating unlike those of other alkali metals which are somewhat stable towards heat.
4LiNO3 → 2Li2O + 4NO2 + O2
2LiCO3 → 2Li2O + CO2
2LiOH → Li2O + H2O
LiOH is a weaker base than hydroxides of other alkali metals.
Unlike elements of group 1, Lithium forms nitride with nitrogen.
3Li + N → Li3N
Lithium halides are have more covalent nature than halides of other members of group 1.

41. Due to appreciable covalent nature, the halides and alkyls of lithium are soluble in organic solvents.
Li + has very high hydration energy and charge/radius ratio, therefore it acts as an excellent reducing agent in solution.
Li +
Li+

42. Small size of atom results in relatively high cohesive properties associated with relatively strong inter-metallic bonding; large atoms usually form weak bonds.

43. Diagonal Relationship
The properties of lithium are quite different from the properties of other alkali metals. On the other hand, it shows greater resemblance with magnesium, which is diagonally opposite element of it. Similarly properties of Beryllium & Boron represent that of Aluminium & Silicon respectively. The main reasons for the anomalous behavior of lithium are -:

44. The Reasons -: (i) The extremely small size of Lithium & its ion
The Reasons -: (i) The extremely small size of Lithium & its ion. (ii) Greater polarizing power of lithium ion ( Li+), due to its small size which result in the covalent character in its compounds. (iii) Least electropositive character and highest ionization energy as compared to other alkali metals. (iv) Non availability of vacant d-orbitals in the valence shell.

45. Some More Examples Examples For Diagonal Relationship
Li and Mg form only normal oxides whereas Na forms peroxide and metals below Na, in addition, forms superoxide.
Li is the only Group 1 element which forms nitride, (Li3N). Mg, as well as other Group 2 elements, also form nitride.
Lithium carbonate, phosphate and fluoride are sparingly soluble in water. The corresponding Group 2 salts are insoluble. (Think lattice and solvation energies).
Both Li and Mg form covalent organometallic compounds. LiMe and MgMe2 (of Grignard reagents) are both valuable synthetic reagents. The other Group 1 and Group 2 analogues are ionic and extremely reactive (and hence difficult to manipulate).

46. THE REACTIVITY OF THESE METALS INCREASES DOWN THE GROUP.
THE ALKALI METALS ARE
HIGHLY REACTIVE. CAUSING
CONTRIBUTING FACTORS ARE LARGE SIZE AND LOW
IONIZATION ENTHALPY.
THE REACTIVITY OF THESE METALS INCREASES DOWN THE GROUP.

47. 1. REACTIVITY TOWARDS AIR -
CHEMICAL
PROPERTIEES
1. REACTIVITY TOWARDS AIR -

48. ALKALI METALS TARNISH IN DRY AIR DUE TO FORMATION OF THEIR OXIDES.
2. THEY BURN IN OXYGEN VIGOURSLY.
3. THEY REACT WITH MOISTURE FORMING HYDROXIDES.

49. EX-
4LI+O LI2O
2NA+O NA2 O2
M+O MO2
OXIDATION STATE- +1

50. LITHUIM IS AN EXEPTION
REACTING DIRECTLY WITH NITROGEN OF AIR TO FORM THE NITRIDE.
DUE TO THEIR HIGH REACTIVITY TO AIR AND WATER THEY ARE KEPT IN
KEROSENE OIL.

51. 2. REACTIVITY TOWARDS WATER-
2M+2H2O M+2OH+H2
THE ALKALI METALS REACT WITH H2O TO FORM HYDROXIDE AND H2.
REACTION WITH H2O IS EXPLOSIVE FOR ALKALI METALS.

52. 3. REACTIVITY TOWARDS HALOGENS-
THE ALKALI METALS READILY REACT VIGOURSLY WITH HALOGENS TO FORM IONIC HALIDES,M+X-.
LITHUIM HALIDES ARE CONVALENT DUE TO POLARISATION.

53. 5. REDUCING NATURE
LITHUIM-STRONGEST
SODIUM-LEAST POWERFUL REDUCING AGENT.
A) M - M (SUBLIMATION ENTHALPY)
B) M=+M+E- (IONISATION ENTHALPY)
C)+M+H2O=M+ (HYDRATION ENTHALPY)

54. SOLUTION IN LIQUID AMMONIA.
THE ALKALI METALS DISSOLVE IN LIQUID AMMONIA GIVING DEEP BLUE SOLUTIONS - CONDUCTING IN NATURE.
THE PRODUCTS OF THIS REACTION ARE -

55. 1.HYDROGEN.
2. AMIDE ION.
THIS REACTION TAKES PLACE AT SLIGHTLY ELEVATED TEMPERATURES OR IN THE PRESENCE OF CATALYST.
WHEN LIQUID AMMONIA IS EXPOSED TO LIGHT OF

56. SPECTRA REGION OF UV LIGHT THERE IS NO OBSERVABLE CHANGE.
WHEN METTALIC SOLUTION ARE SO EXPOSED,REACTION OCCURS.THIS REACTION IN PRESENCE OF UV IS COMPLETE PHOTOCHEMICAL REACTION.
REACTION-

57. THE BLUE COLUR IS DUE TO THE AMMONIATED ELECTRON
+M + E- + NH MNH2 +1/2H2
THE BLUE COLUR IS DUE TO THE AMMONIATED ELECTRON
WHICH ABSORBS ENERGY IN VISIBLE REGION OF LIGHT IMPARTING BLUE COLOUR TO SOLUTION. THE SOLUTION IS PARAMAGNETIC.

58. IN CONCENTRATED SOLUTION THE BLUE COLOUR CHANGES TO BRONZE COLOUR AND BECOME DIMAGNETIC.

59. 1. LITHUIM IS USED TO MAKE ALLOYS
1.LITHUIM IS USED TO MAKE ALLOYS.EX-WITH LEAD TO MAKE WHITE METAL,BEARINGS FOR MOTORS ENGINES,WITH ALUMINIUM TO MAKE AIRCRAFTS PARTS AND WITH MAGNESIUM TO MAKE ARMOUR PLATES.
LITHUIM IS ALSO USED TO MAKE ELECTROCHEMICAL CELLS.
USES

60. 2. SODIUM IS USED TO MAKE A Na or Pb ALLOY NEEDED TO MAKE PbMe4
2. SODIUM IS USED TO MAKE A Na or Pb ALLOY NEEDED TO MAKE PbMe4. used as anti-knock additives to petrol.
Liquid sodium metal is used as a coolant in fast breeder nuclear REACTORS.
3. POTASSIUM HAS A VITAL ROLE IN BIOLOGICAL SYSTEMS.

61. POTASSIUM CHLORIDE IS USED AS A FERTILISER.
POTASSIUM HYDROXIDE IS USED AS AN EXCELLENT ABSORBENT OF CARBON DIOXIDE.
4.CAESIUM IS USED IN DEVISING PHOTOELECTRIC CELLS.