THE ALKALINE EARTH GROUP OF IONS
Ba2+ , Sr2+ , Ca2+ and Mg2+
Uses : Mg in alloys and increase weight of aircraft structure
Ca in biological system for bone growth
Sr radioactive isotope
Ba its sulphate is opaque to X-rays
Ba2+ , Sr2+ and Ca2+
• Oxides and Hydroxides
Reaction of the oxides with water yields the metal hydroxides which
increase in solubility with increased atomic weight of the metal. Thus while
Mg(OH)2 is slightly soluble, Ba(OH)2 is freely soluble.
• Carbonates
The carbonates of Ca2+ , Sr2+ and Ba2+ are precipitated on the addition of
alkali carbonate or ammonium carbonate. Mg2+ ion under such condition is
partially precipitated as basic magnesium carbonate. However, in ammonium
carbonate solution having the CO32- ion concentration regulated by the presence
of both NH4+ ions and free NH3 in moderately high concentrations, Mg2+ ion
remains in solution. This is according to the following equilibrium:
NH3 + HCO3-
NH4+ + CO32-
• Chromates
The molar solubility of alkaline earth chromates is increased in the
following order
BaCrO4 < SrCrO4 < CaCrO4
The molar solubility of CaCrO4 is 200 times greater than that of SrCrO4
and the latter is 600 times greater than that of BaCrO4.
• Sulfates
The molar solubility of alkaline earth sulfates is increased in the order
BaSO4 < SrSO4 < CaSO4
-1-
• Oxalates
The molar solubilities of SrC2O4 and BaC2O4 are only ten times greater
than that for CaC2O4. Hence the separation of these cations as oxalates is
impractical.
The oxalates are freely soluble in strong acids. SrC2O4 and BaC2O4 are
appreciably soluble in acetic acid while CaC2O4 is insoluble.
• Phosphates
The alkaline earth cations are precipitated in the form of MHPO4 on
addition of Na2HPO4 to the neutral solution. In strongly alkaline solution
M3(PO4)2 is precipitated. In acidic medium the phosphates are soluble due to
the formation of H2PO4- ion or H3PO4.
• Complexes
The stability of the complexes increases with increased charge density,
thus the complexes of Ca2+ ion are the most stable. Ca2+ ion forms a moderately
stable chelate with triethanolamine and a very stable one with ethylenediamine
tetra-acetic acid (EDTA). Ba2+ and Sr2+ ions, being of low charge density, their
complexes are not common.
MAGNESIUM ION (Mg2+)
Salts of analytical importance
• Magnesium hydroxide
It is precipitated from aqueous solution of Mg2+ ion by the addition of
alkali hydroxide or ammonium hydroxide. Mg(OH)2 is not amphoteric and does
not form ammine complex. It is precipitated neither from NH4OH-NH4Cl
buffered solution nor from NH4Cl and (NH4)2CO3 buffered medium. NH4+
represses ammonia ionization to give low [OH-] with the result that the
solubility product for Mg(OH)2 is not exceeded.
NH4Cl → NH4+ + ClNH4+ + OH-
NH4OH
-2-
• Magnesium phosphate
Addition of HPO42- ions to Mg2+ ions precipitates a gelatinous precipitate
of Mg HPO4.7H2O in neutral aqueous medium or Mg3(PO4)2 in strongly alkaline
medium or Mg(NH4)PO4. 6H2O in an aqueous solutions containing NH4+ and
NH3.
Mg2+ + HPO42- + 7H2O = MgHPO4.7H2O
3Mg2+ + 2HPO42- + 2OH- = (Mg)3 (PO4)2 + 2H2O
Mg2+ + NH4+ + PO43- + 6H2O = MgNH4PO4.6H2O
-3-
THE ALKALI METAL GROUP OF IONS
(K+ and Na+); and NH4+ ion
• The alkali metals resemble one another in their physical and chemical
properties. The metals have a single valence electron which is easily lost,
therefore they are active reducing agents. They are rapidly attacked by
oxygen or moisture; hence, they cannot exist as such when exposed to
atmosphere. They form hydrides (e.g. Na+H-) which behave as typical ionic
compounds.
• Ammonium ion, on the other hand, formed by the addition of a proton to
the NH3 molecule carries one positive charge. The solubility of its salts
resembles that of the K+ ions as might be predicted from the similarity in
the charge and radius. Therefore, ammonium ion is included in the same
analytical group of alkali metals. Ammonium ion, however, differs from
alkali metal ions in (a) it forms a relatively weak base and (b) its salts
decompose on heating.
heat
NH4Cl → NH3 + HCl
NH4+ + NO3- = N2O + 2H2O
Sodium ion (Na+) and potassium ion (K+)
• The alkali metal salts of strong acids give neutral solutions. The salts of
weak acids form alkaline solutions by hydrolysis of the anions, which
capture proton from water.
NaAc → Na+ + Ac+H2O
HAc + OH-
-4-
K+ + ClO4- = KClO4
white ppt.
K+ + H2C4H4O6 = KHC4H4O6 + H+
white ppt.
2K+ + PtCl62- = K2PtCl6
yellow ppt.
K+ + Na+ + Co(NO2)63- = K2Na Co(NO2)6 or K3Co(NO2)6
yellow ppt (used as a spot test)
Na+ + HMg(UO2)3 Ac9 + 6H2O = Na Mg(UO2)3 Ac9.6H2O + H+
yellow ppt (used as a spot test)
• Other properties
(1) The oxides dissolve in H2O to give strong alkalis.
(2) The carbonates are water-soluble giving CO2 with acids but do not give
CO2 on heating.
(3) Halides give, HX or X2 with conc. H2SO4.
(4) Nitrates give HNO3 with conc. H2SO4 and give O2 (not NO2) when
heated.
2KNO3 → KNO2 + O2
NaNO3 + H2SO4 → NaHSO4 + HNO3
(5) Sulphites give SO2 on warming with acids.
(6) Thiosulphates :
• Prepared on boiling sulphite with sulphur.
• Deposit s on treating with H+
• Used in iodine titration
• The presence of NH4+ in solution may be detected through addition of
excess strong base to convert the ion to molecular ammonia. The latter can
be detected by (a) its odour, (b) alkaline character to moistened litmus, (c)
white clouds of NH4Cl with HCl; (d) blackening a paper of mercurous
nitrate; and (e) use of Nessler's reagent - (if the quantity is minute) – to
produce a yellow-brown colloidal suspension or precipitate.
The last detection can be represented by the following equation.
2Hg I42- + 3OH- + NH3 = NH2Hg2I3 + 5I- + 2H2O
-5-