Q.Why is Li2CO3 decomposed at a lower temperature whereas Na2CO3 at higher temperature?

Anomalous Behaviour of Lithium

You already know that all alkali metals sit together in Group 1 — sodium, potassium, rubidium, caesium, francium. They are famously reactive, soft, and form strongly ionic compounds. Lithium sits right at the top of that column, and you might expect it to behave just like the others. It doesn't. In several important ways, lithium acts more like magnesium (a Group 2 element) than like its own group members. That is the anomaly.

Why does this happen?

Two properties of lithium are responsible for almost everything unusual about it.

First, its size. Lithium is the smallest metal atom in the entire periodic table. Its atomic radius is about 152 pm, while sodium is already 186 pm. That tiny size means its outermost electron is held very tightly by the nucleus — the ionisation enthalpy of lithium is the highest among alkali metals.

Second, its polarising power. When lithium forms a positive ion, Li⁺, that ion is extremely small (only about 76 pm). A small, highly charged cation distorts the electron cloud of any anion it meets — this is polarisation. High polarising power pushes the bonding away from pure ionic character and toward covalent character. For lithium, this effect is strong enough to change the chemistry.

The precise statement

The anomalous behaviour of lithium means that lithium differs significantly from the rest of the alkali metals in several of its chemical properties, and in many of these it resembles magnesium instead. The key differences are:

1. Lithium forms a covalent, unstable carbonate. All other alkali metal carbonates (Na₂CO₃, K₂CO₃, etc.) are ionic, thermally stable, and do not decompose on heating. Lithium carbonate, Li₂CO₃, is much less ionic — it is partially covalent — and it decomposes on heating into lithium oxide and carbon dioxide:

This is exactly what magnesium carbonate does. Sodium carbonate, by contrast, can be heated red-hot without decomposing.

2. Lithium reacts with nitrogen to give a nitride. No other alkali metal does this. When lithium is heated in air, it combines directly with nitrogen to form lithium nitride:

Magnesium does the same thing — it forms Mg₃N₂. Sodium and potassium simply do not react with nitrogen at all.

3. Lithium forms a stable oxide (Li₂O) rather than a peroxide or superoxide. Sodium gives Na₂O₂ (peroxide) when burned in air; potassium, rubidium, and caesium give superoxides (KO₂, RbO₂, CsO₂). Lithium, because of its small cation size, can only stabilise the simple oxide Li₂O. Again, magnesium gives MgO.

4. Lithium hydroxide is a weaker base. LiOH is moderately strong but much less basic than NaOH or KOH. It decomposes on heating to Li₂O, while the other alkali metal hydroxides do not decompose.

5. Lithium salts are often hydrated and soluble in organic solvents. LiCl is deliquescent and soluble in alcohol; NaCl and KCl are not. The small Li⁺ ion strongly attracts water molecules, forming hydrates like LiCl·3H₂O. …