Concept understanding — Anomalous Behaviour of Lithium
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.
Note
The combination of small size and high polarising power is often called the diagonal relationship — lithium resembles magnesium, which sits diagonally below and to the right in the periodic table.
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:
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:
Li2CO3ΔLi2O+CO2
This is exactly what magnesium carbonate does. Sodium carbonate, by contrast, can be heated red-hot without decomposing.
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:
6Li+N2→2Li3N
Magnesium does the same thing — it forms Mg₃N₂. Sodium and potassium simply do not react with nitrogen at all.
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.
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.
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. …
Lithium's unusually small ionic size and high charge density make it behave differently from the rest of Group 1, a pattern often called the anomalous behaviour of lithium.
Lithium's unusually small size and high charge density make it behave differently from the rest of Group 1 (the "anomalous behaviour of lithium").
Because Li+ is by far the smallest alkali metal cation, it has an exceptionally high charge density and polarising power, causing it to differ from the rest of Group 1:
Hardness and melting/boiling point: Lithium is the hardest of the alkali metals and has the highest melting/boiling points in the group. Reason: its small atomic size gives stronger metallic bonding compared to the larger, softer atoms of Na, K, Rb, Cs.
Nature and hydration of salts: LiCl shows much greater covalent character than the chlorides of other alkali metals, is deliquescent, crystallises as a hydrate (LiCl.2H2O), and is soluble in organic solvents like ethanol. The chlorides of other alkali metals are typically anhydrous and largely ionic. Reason: Li+'s very small size polarises the large Cl- ion (per Fajan's rules), imparting covalent character, and its high charge density gives a strong tendency to hydrate.
Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.
CBSE 2025Set ANNUAL2 marks
Q.Lithium salts are mostly hydrated. Why?
›Reveal solutionSolution
Lithium salts are mostly hydrated because the very small, highly charge-dense Li⁺ ion has a strong ion-dipole attraction for water molecules.
Among the alkali metals (Group 1), the ionic radius increases down the group (Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺). Since all these ions carry the same +1 charge, the charge density (charge/volume) is highest for Li⁺, the smallest ion.
A high charge density means the ion exerts a strong electric field around itself. Water is a polar molecule (dipole), and Li⁺'s strong field pulls the oxygen (partially negative) end of several water molecules towards it, forming a tightly bound hydration shell. The hydration enthalpy of Li⁺ is therefore much larger (more negative) than that of the heavier alkali metal ions.
Li+'s very small ionic size gives it an unusually high charge density, so it attracts and holds water molecules of hydration far more strongly than the other alkali metal ions.
Within Group 1, ionic size increases down the group (Li+<Na+<K+<Rb+<Cs+). Since all alkali metal ions carry the same +1 charge, the smallest ion — Li+ — has the highest charge density (charge concentrated over the smallest volume) and therefore the strongest electrostatic field around it.
This strong field polarises and strongly attracts the polar water molecules, so Li+ has the highest hydration enthalpy of all the alkali metal ions. Consequently:
Lithium salts are usually obtained and exist as hydrates, e.g. LiCl⋅2H2O, LiClO4⋅3H2O. …
Lithium's small ionic size gives Li+ an unusually high charge density, so it hydrates strongly — most lithium salts are found as hydrates (e.g. LiCl.2H2O).
Lithium is the smallest alkali metal, and its cation Li+ (ionic radius ~76 pm) is by far the smallest of the Group 1 cations. Charge density (charge/size) of an ion determines how strongly it can polarise and attract the negative (oxygen) end of surrounding water dipoles.
Because Li+ has a very small radius carrying a full unit positive charge, its charge density and hence its polarising power are much higher than the other alkali metal ions (Na+, K+, ...). This lets Li+ attract water molecules strongly around itself, releasing a large hydration enthalpy in the process. As a result:
Li+ in aqueous solution and in many of its salts is surrounded by a firmly held shell of water molecules.
Most lithium salts crystallise as hydrates (e.g. LiCl.2H2O, LiClO4.3H2O), unlike the corresponding sodium/potassium salts which are often anhydrous. …
Li+ is very small with a high charge density, so it strongly attracts water molecules; hence lithium salts are mostly hydrated.
The tendency of a salt to be hydrated depends on the polarising power (charge density) of its cation. Charge density = charge / size; the smaller the ion (for the same charge), the greater its charge density and the more strongly it attracts polar water molecules.
Among the alkali metal ions (Li+, Na+, K+, Rb+, Cs+), the lithium ion Li+ is the smallest. Therefore:
Li+ has the highest charge density and the greatest hydrating (and polarising) power of all the alkali metal cations.
It attracts water molecules very strongly, so a large amount of hydration energy is released and the water becomes firmly attached.
…
The small size of Li⁺ gives it a very high charge-to-size ratio, so it has a strong tendency to attract and hold water molecules, forming hydrated crystalline salts.
Lithium is the first member of Group 1 (alkali metals) and has anomalous behaviour compared to the rest of its group because of its very small ionic size.
The ionic radius of Li⁺ is the smallest among all the alkali metal ions.
A small ionic radius means a high charge density (charge concentrated over a small volume), which gives the Li⁺ ion a strong electric field around it.
This strong electric field polarises and strongly attracts the polar water molecules (through ion–dipole interaction), causing several water molecules to cluster tightly around each Li⁺ ion. …
Q.List any two points of difference between Lithium and other alkali metals. Give reasons.
›Reveal solutionSolution
Lithium's unusually small size and high charge density make it behave differently from the rest of Group 1 (the "anomalous behaviour of lithium").
Because Li+ is by far the smallest alkali metal cation, it has an exceptionally high charge density and polarising power, causing it to differ from the rest of Group 1:
Hardness and melting/boiling point: Lithium is the hardest of the alkali metals and has the highest melting/boiling points in the group. Reason: its small atomic size gives stronger metallic bonding compared to the larger, softer atoms of Na, K, Rb, Cs.
Nature and hydration of salts: LiCl shows much greater covalent character than the chlorides of other alkali metals, is deliquescent, crystallises as a hydrate (LiCl.2H2O), and is soluble in organic solvents like ethanol. The chlorides of other alkali metals are typically anhydrous and largely ionic. Reason: Li+'s very small size polarises the large Cl- ion (per Fajan's rules), imparting covalent character, and its high charge density gives a strong tendency to hydrate.