Oxides and hydroxides
All the alkali metals ,their oxides, peroxide and superoxide readily dissolve in water to produce corresponding hydroxides which are strong alkalies.
2 Na + 2 H2O ————-> 2 NaOH + H2
Na2O + H2O ————-> 2 NaOH
Na2O2 + 2 H2O ————-> 2 NaOH + H2O2
2KO2 + 2 H2O ————-> 2 KOH + H2O2 + O2
Peroxides and superoxides also act as oxidising agents since they react with water forming H2O2 and O2.
The hydroxides of all the alkali metals are white crystalline solids. They are strongest of all bases and readily dissolve in water with the evolution of much heat. A number of hydrates of the heavier alkali metal hydroxides have been prepared from their aqueous solution but little is known about their structure.
1) Basic strength : The basic strength of these hydroxide increases as we move down the group from Li to Cs. The hydroxides of alkali metals behave as strong bases due to their low ionisation enthalpies. The M—O bond in M—O—H can easily break giving M+ and OH‾ ions.
As we move down the group ,the ionisation enthalpy decreases. As a result , the M–OH the bond is more and more easily cleaved and hence the basic strength increases down the group from LiOH to CsOH.
2) Solubility and stability: All these hydroxides are highly soluble in water and thermally stable except Lithium hydroxide.
2 LiOH ——–> Li2O + H2O
3) Formation of salts with acid : Alkali metal hydroxides being strongly basic react with all acids forming salts.
NaOH + HCl ———> NaCl + H2O
The salts are colourless ionic solids which are soluble in water.
The alkali metals combine directly with halogens under appropriate conditions forming halides of the general formula, MX. These halides can also be prepared by the action of aqueous halogen acids on metal oxides ,hydroxides or carbonates.
M2O + 2 HX —–> 2 HX + H2O
MOH + HX ——> MX + H2O
M2CO3 + 2 HX ———> 2 MX + CO2 + H2O
where M= Li, Na, K, Rb, Cs
X = F, Cl, Br, I
All these halides are colourless ,high melting crystalline solids having high negative enthalpy of formation.
The ΔHf ° values for fluorides become less and less negative as we move down the group while for chlorides, bromides and iodides these values become more and more negative.
Fluorides are the most stable while iodides are the least stable.
a) Polarization effect
When a cation approaches an anion, the electron cloud of the anion is attracted towards the cation and hence get distorted. This effect is called polarization. The power of the cation to polarize the anion is called its polarising power and the tendency of the anion to get polarized is called its polarizability.
The greater the polarization produced more is the concentration of electrons between the two atoms thereby decreasing the ionic character or increasing the covalent character of the bond.
The covalent character of any compounds, depends upon:
1)Size of the cation : smaller the cation, greater is its polarising power and hence larger is the covalent character. The covalent character decreases as the size of the cation increases.
LiCl > NaCl> KCl > RbCl> CsCl.
LiCl is more covalent than KCl.
2) Size of the anion : Larger the anion, greater is its polarisability. Covalent character of lithium halides is in the order: LiI > LiBr > LiCl> LiF
Dipole moment of LiI is much less than the theoretically expected value of 11.5 D if it were 100 % ionic.
3)Charge on the ion : Greater the charge on the cation ,greater is its polarising power and hence larger is the covalent character. the covalent character of some of the halides increases in the order :
Na+ Cl‾ < Mg2+ ( Cl‾)2 < Al3+ ( Cl‾)3
Greater the charge on the anion, more easily it gets polarised thereby imparting more covalent character to the compounds formed.
Covalent character increases in the order: NaCl < Na2SO4 < Na3PO4 as the size of the anion increases and hence larger is the covalent character. The covalent character decreases as the size of the anion decreases.
4) Electronic configuration of the cation: If two cations have the same charge and size, the one with a pseudo noble gas configuration i.e. having 18 electrons in the outermost shell has greater polarising power then a cation with noble gas configuration i.e. having 8 electrons in the outermost shell.
b) lattice enthalpies: lattice enthalpy is defined as the amount of energy required to separate out one mole solid ionic compound into its gaseous ion.Greater the lattice enthalpy, higher is the melting point of the alkali metal halide and lower if its solubility in water.
c) Hydration enthalpy : It is the amount of enthalpy released when one more of gaseous ions combine with water to form hydrated ions.
M+ (g) + aq ——–> M+ (aq) + hydration enthalpy
X‾(g) + aq ——–> X‾ (aq) + hydration enthalpy
The hydration enthalpy of the ions ,greater is the solubility of the salt in water.
The extent of hydration depend upon the size of the ion. Smaller the size of the ion, mor highly it is hydrated and hence greater is its hydrated ionic radius and less is its ionic mobility.
1) A balance between lattice enthalpy and hydration enthalpy determines the ultimate solubility of a compound in water.
For Ex: LiF is almost insoluble in water due to its high lattice energy. The low solubility of CsI is due to smaller hydration energy of the two ions.
2) The solubility of most of the alkali metal halides except those of fluoride decreases on descending the group since the decrease in hydration enthalpy is more than the corresponding decrease in the lattice enthalpy.
3) Because of the small size and higher electronegativity, Lithium halides except LiF are predominantly covalent and hence are soluble in organic solvents such as alcohol, acetone, ethyl acetate. LiCl is also soluble in pyridine.NaCl being ionic is insoluble in organic solvents.
4) Due to high hydration enthalpy of Li+ ion, Lithium halides are soluble in water except LiF which is sparingly soluble due to its high lattice enthalpy. As we move down the group ,the solubility of alkali metal fluorides increases regularly as we move from LiF to CsF since the decrease in lattice enthalpy more than compensate the decrease in hydration enthalpy.
5) For the same alkali metals, the melting point decrease in the order : Fluoride > Chloride > bromide> iodide.
As the lattice energy decrease , energy required to break the lattice decrease and hence the melting point of sodium halides decreases from NaF —-> NaCl ——> NaBr —–> NaI.
6) 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.
Reason: The low melting point of LiCl as compared to that of NaCl is probably because LiCl is covalent in nature while NaCl is ionic.
Salts of oxoacids
Since the alkali metals are highly electropositive, therefore , their hydroxides are very strong bases and hence they forms salts with all oxoacids.
Oxoacids are those in which the acidic proton is on a hydroxyl group.
For Ex: H2CO3, H3PO4, H2SO4, HNO3, HNO3.
They are generally soluble in water and stable towards heat.
Carbonates: The carbonates of alkali metals are quite stable towards heat. Li2CO3 is less stable and decompose readily on heating. This is due to the reason that small lithium cation polarizes the oxygen atom of the nearby larger carbonate ion. As a result, the C-O bond weakens and Li-O bond strengthens thereby favouring the formation of Li2O and CO2 . It is the higher lattice energy of Li2O over that of Li2CO3 which favours the decomposition of Li2CO3.
As the size of the cation increases from Na+ to Cs+ , the lattice energy of their corresponding oxides decreases and hence the stability of the carbonate increases from Na2CO3 to Cs2CO3.
Bicarbonate: Being strongly basic, alkali metals also form solid bicarbonates. No other metal form solid bicarbonates, though NH4HCO3 also exist as a solid. Lithium, however does not form solid bicarbonates though it does not exist in solution. All the bicarbonates on gentle heating undergo decomposition to form carbonate with the evolution of carbon dioxide.
2 MHCO3 ——–> M2CO3 + CO2 + H2O
As the basic character of metal hydroxide increases from LiOH to CsOH or the electropositive character of the metal increases from Li to Cs, the stability of carbonates and bicarbonates increases.
All the carbonates and bicarbonates are soluble in water and their solubility increases rapidly on descending the group. This is due to the reason that their lattice energy decrease more rapidly than their hydration energy on moving down the group.