`color{green}(★)` On combustion in excess of air, lithium forms mainly the oxide, `color{red}(Li_2O)` (plus some peroxide `color{red}(Li_2O_2)` ) sodium forms the peroxide `color{red}(Na_2O_2)` (and some superoxide `color{red}(Na O_2)` ) whilst potassium, rubidium and caesium form the superoxides, `color{red}(MO_2)` .

`color{green}(★)` Under appropriate conditions pure compounds `color{red}(M_2O , M_2O_2)` and `color{red}(MO_2)` may be prepared.

`color{green}(★)` The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilisation of large anions by larger cations through lattice energy effects. These oxides are easily hydrolysed by water to form the hydroxides according to the following reactions :

`color{red}(M_2O + H_2O → 2M^(+) + 2OH^(-))`

`color{red}(M_2O_2 + 2H_2O → 2M^(+) + 2OH^(-) + H_2O_2)`

`color{red}(2MO_2+2H_2O → 2M^(+) + 2OH^(-) + H_2O_2 + O_2)`

`color{green}(★)` The oxides and the peroxides are colourless when pure, but the superoxides are yellow or orange in colour.

`color{green}(★)` The superoxides are also paramagnetic. Sodium peroxide is widely used as an oxidising agent in inorganic chemistry.

`color{green}(★)`The hydroxides which are obtained by the reaction of the oxides with water are all white crystalline solids.

`color{green}(★)` The alkali metal hydroxides are the strongest of all bases and dissolve freely in water with evolution of much heat on account of intense hydration.

`color{green}(★)` The alkali metal halides , `color{red}(MX , ( X =F , Cl , Br , I))` are all high melting, colourless crystalline solids.

`color{green}(★)` They can be prepared by the reaction of the appropriate oxide, hydroxide or carbonate with aqueous hydrohalic acid (`color{red}(HX)`).

`color{green}(★)` All of these halides have high negative enthalpies of formation; the `color{red}(△_f H^(⊖))` values for fluorides become less negative as we go down the group whilst the reverse is true for `color{red}(△_fH^(⊖))` for chlorides bromides and iodides. For a given metal `color{red}(△_fH^(⊖))` always becomes less negative from fluoride to iodide.

`color{green}(★)` The melting and boiling points always follow the trend : fluoride > chloride > bromide > iodide. All these halides are soluble in water.

`color{green}(•)` The low solubility of `color{red}(LiF)` in water is due to its high lattice enthalpy whereas the low solubility of `color{red}(CsI)` is due to smaller hydration enthalpy of its two ions.

`color{green}(•)` Other halides of lithium are soluble in ethanol, acetone and ethylacetate; `color{red}(LiCl)` is soluble in pyridine also.

`color{green}(★)` Oxo-acids are those in which the acidic proton is on a hydroxyl group with an oxo group attached to the same atom e.g., carbonic acid, `color{red}(H_2CO_3 \ \ \ ( OC (OH)_2)) `; sulphuric acid, `color{red}(H_2SO_4 (O_2S(OH)_2 ))`.

`color{green}(★)` The alkali metals form salts with all the oxo-acids.

`color{green}(★)` They are generally soluble in water and thermally stable.

`color{green}(★)` Their carbonates `color{red}((M_2CO_3))` and in most cases the hydrogencarbonates `color{red}((MHCO_3))` also are highly stable to heat.

`color{green}(★)` As the electropositive character increases down the group, the stability of the carbonates and hydorgencarbonates increases.

`color{green}(•)` Lithium carbonate is not so stable to heat; lithium being very small in size polarises a large `color{red}(CO_3^(2-))` ion leading to the formation of more stable `color{red}(Li_2O)` and `color{red}(CO_2)` .

`color{green}(•)` Its hydrogen carbonate does not exist as a solid.