Alkali Metals (Group I elements of modern periodic table):
Metallic radius (pm)
Ionic radius (M+/pm)
Density/g cm–3 (at 298K)
E°(V) at 298K for
M+(aq) + e–→ M(s)
*ppm (parts per million)
** percentage by weight
Physical Properties of Alkali Metals:
Hydroxides of Alkali Metals:
a)All the alkali metals, their oxides, peroxides and superoxides readily dissolve in water to produce corresponding hydroxides which are strong alkalies.
b) The basic strength of these hydroxides increases as we move down the group Li to Cs.
c) All these hydroxides are highly soluble in water and thermally stable except lithium hydroxide.
d) Alkali metals hydroxides being strongly basic react with all acids forming salts.
Halides of Alkali metals:
M2O + 2HX → 2MX + H2O
MOH + HX → MX + H2O
M2CO3 + 2HX → 2MX + CO2 + H2O (M = Li, Na, K, Rb or Cs)
(X = F, Cl, Br or I)
a) Standard enthalpies of formation in (kJ/mol-1)
b) Covalent Character:.
c) Lattice Energies: Amount of energy required to separate one mole of solid ionic compound into its gaseous ions.
Greater the lattice energy, higher is the melting point of the alkali metals halide and lower is its solubility in water
d) Hydration Energy: Amount of energy released when one mole of gaseous ions combine with water to form hydrated ions.
Higher the hydration energy of the ions greater is the solubility of the compound in water.
The solubility of the most of alkali metal halides except those of fluorides decreases on descending the group since the decrease in hydration energy is more than the corresponding decrease in the lattice energy.
Due to high hydration energy of Li+ ion, Lithium halides are soluble in water except LiF which is sparingly soluble due to its high lattice energy.
For the same alkali metal the melting point decreases in the order
fluoride > chloride > bromide > iodide
For the same halide ion, the melting point of lithium halides are lower than those of the corresponding sodium halides and thereafter they decrease as we move down the group from Na to Cs.
The low melting point of LiCl (887 K) as compared to NaCl is probably because LiCl is covalent in nature and NaCl is ionic.
Anomalous Behavior of Lithium and diagonal relationship with Magnesium:
Li has anomalous properties due to
Lithium show diagonal relationship with magnesium because both elements have almost same polarizing power.
Lithium nitrate, on heating, decomposes to give lithium oxide, Li2O whereas other alkali metals nitrate decomposes to give the corresponding nitrite.
Sodium Hydroxide (NaOH):
Sodium Carbonate (Washing soda) (Na2CO3):
Carbon dioxide gas is bubbled through a brine solution saturated with ammonia and itresults in the formation of sodium hydrogen carbonate.
Sodium hydrogen carbonate so formed precipitates out because of the common ion effect caused due to the presence of excess of NaCl. The precipitated NaHCO3 is filtered off and then ignited to get Na2CO3.
2NaHCO3 → Na2CO3 + CO2 + H2O
1. The aqueous solution absorbs CO2 yielding sparingly soluble sodium bicarbonate.
2. dissolves in acids with an effervescence of carbondioxide and is causticised by lime to give caustic soda.
3. Fusion with silica, sodium carbonate yields sodium silicate.
4. Hydrolysis – being a salt of a strong base (NaOH) and weak acid (H2CO3), when dissolved in water sodium carbonate. Undergoes hydrolysis to form an alkaline solution
Alkali Earth Metals (Group II elements of modern periodic table):
Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra).
1. Alkali earth metals are almost similar in properties to the corresponding alkali metals.
2. Atomic and ionic radii
The atomic radii as well as ionic radii of the members of the family are smaller than the corresponding members of alkali metals.
3. Ionization energy: The alkaline earth metals owing to their large size of atoms have fairly low values of ionization energies as compared to the p – block elements. However with in the group, the ionization energy decreases as the atomic number increases. It is because of increase in atomic size due to addition of new shells and increase in the magnitude of screening effect of the electrons in inner shells. Because their (IE)1 is larger than that of their alkali metal neighbours, the group IIA metals trend to the some what less reactive than alkali metals. The general reactivity trend is Ba > Sr > Ca > Mg > Be.
4. Oxidation state: The alkaline earth metal have two electrons in their valence shell and by losing these electrons, these atoms acquire the stable noble gas configuration. Thus, unlike alkali metals, the alkaline earth metals exhibit +2 oxidation state in their compounds.
M → M+2 + 2e-
5. Characteristic flame colouration:
Carmine – red
Difference between alkali metals and alkali earth metals:
Alkaline earth metals
Two electrons are present in the valency shall. The configuration is ns2 (bivalent)
One electron is present in the valency shell. The configuration is ns1 (monovalent) more electropositive
Weak bases, less soluble and decompose on heating.
Strong bases, highly soluble and stable towards heat.
These are not known in free state. Exist only in solution.
These are known in solid state.
Insoluble in water. Decompose on heating.
Soluble in water. Do not decompose on heating (LiCO3 is an exception)
Action of nitrogen
Directly combine with nitrogen and form nitrides
Do not directly combine with nitrogen except lithium
Action of carbon
Directly combine with carbon and form carbides
Do not directly combine with carbon
Decompose on heating evolving a mixture of NO2 and oxygen
Decompose on heating evolving only oxygen
Solubility of salts
Sulphates, phosphates fluorides, chromates, oxalates etc are insoluble in water
Sulphates, phosphates, fluorides, chromates, oxides etc are soluble in water.
Comparatively harder. High melting points. Diamagnetic.
Soft, low melting points paramagnetic.
Hydration of compounds
The compounds are extensively hydrated. MgCl2.6H2O, CaCl2.6H2O, BaCl2.2H2O are hydrated chlorides.
The compounds are less hydrated. NaCl, KCl, RbCl form non – hydrated chlorides
Weaker as ionization potential values are high and oxidation potential values are low.
Stronger as ionization potential values are low and oxidation potential values are high.
Chemical Properties of Alkali Earth Metals:
1. Reaction with water :
2. Formation of oxides and nitrides
3. Formation of Nitrides
4. Reaction with hydrogen:
M + H2 + Δ → MH2
Both BeH2 and MgH2 are covalent compounds having polymeric structures in which H – atoms between beryllium atoms are held together by three
centre – two electron (3C - 2e) bonds as shown below:
5. Reaction with carbon – (Formation of carbides)
When BeO is heated with carbon at 2175 – 2275 K a brick red coloured carbide of the formula Be2C is formed
It is a covalent compound and react water forming methane.
Be2C + 4H2O → 2Be (OH)2 + CH4
6. Reaction with Ammonia:
Like alkali metal, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solution from which ammoniates [ M (NH3)6 ]2+ can be recovered.
Anamolous Behaviour of Beryllium:
Diagonal relationship of Be with Al:
Calcium Carbonate (CaCO3):
It occurs in nature as marble, limestone, chalk, coral, calcite, etc. It is prepared as a white powder, known as precipitated chalk, by dissolving marble or limestone in hydrochloric acid and removing iron and aluminium present by precipitating with NH3, and then adding ammonium carbonate to the solution; the precipitate is filtered, washed and dried.
CaCl2 + (NH4)2CO3 →CaCO3 + 2NH4Cl
It dissolves in water containing CO2, forming Ca(HCO3)2 but is precipitated from solution by boiling.
CaCO3 + H2O + CO2 ↔ Ca(HCO3)2
Plaster of Paris, CaSO4.1/2 H2O or (CaSO4)2.H2O:
It occurs in nature as gypsum and the anhydrous salt as anhydride. It is prepared by precipitating a solution of calcium chloride or nitrate with dilute sulphuric acid.
The effect of heat on gypsum or the dihydrate presents a review of interesting changes. On heating the monoclinic gypsum is first converted into orthorhombic form without loss of water. When the temperature reaches 120°C, the hemihydrate or plaster of paris is the product. The latter losses water, becomes anhydrous above 200°C and finally above 400°C, it decomposes into calcium oxide.
2CaSO4 → 2CaO + 2SO2↑ + O2↑
The addition of common salt accelerates the rate of setting, while a little borax or alum reduces it. The setting of plaster of paris is believed to be due to rehydration and its reconversion into gypsum.
2CaSO4. 1/2 H2O + 3H2O → 2CaSO4. 2H2O
Plaster of Paris Gypsum
Industrial uses of lime and Limestone
Uses of lime
Uses of Slaked lime [Ca(OH)2]
Uses of lime stone (CaCO3)