2_2 Anomalous Lithium and Beryllium and the Diagonal Relationship

The first element in each main group tends to break the pattern set by the rest of the group, and the s-block shows this clearly. Lithium is a reluctant alkali metal and beryllium an unusual alkaline earth metal. In many ways each behaves more like the element one place down and to the right of it, magnesium for lithium and aluminium for beryllium. This diagonal relationship has a single cause, the very high charge density of the small

\mathrm{Li^+}

and

\mathrm{Be^{2+}}

ions.

The Root Cause: Charge Density and Polarizing Power

Charge density is roughly an ion's charge divided by its size.

\mathrm{Li^+}

(radius about 60 pm) has only one filled shell under it, and

\mathrm{Be^{2+}}

(about 31 pm) carries a double charge on an even smaller ion. A small, highly charged cation pulls the electron cloud of a nearby anion towards itself. This is called polarization, and it shares some of the anion's electrons with the cation, so the bond gains covalent character. The same compact size also makes lithium and beryllium less electropositive than their heavier neighbours, with higher ionization energies and electronegativities.

Diagonal pairs have similar charge density

Ion	Li⁺	Mg²⁺	Be²⁺	Al³⁺
Charge	+1	+2	+2	+3
Ionic radius (pm)	60	65	31	50
Behaves like	Mg²⁺	Li⁺	Al³⁺	Be²⁺

Moving one step right across a period adds charge and shrinks the ion, which raises its charge density. Moving one step down adds a shell, which lowers it. The two changes roughly cancel along a diagonal, so

\mathrm{Li^+}

and

\mathrm{Mg^{2+}}

distort anions to a similar extent, and so do

\mathrm{Be^{2+}}

and

\mathrm{Al^{3+}}

.

How Lithium Differs from the Other Alkali Metals

• It is harder, has a higher melting point and reacts only steadily with water, while sodium and potassium react violently.
• Burnt in air it gives only the normal oxide

  \mathrm{Li_2O}

  . Sodium gives the peroxide and potassium the superoxide.
• It is the only alkali metal that combines directly with nitrogen, forming lithium nitride:

  \mathrm{6Li + N_2 \rightarrow 2Li_3N}

  .
• LiF,

  \mathrm{Li_2CO_3}

  and

  \mathrm{Li_3PO_4}

  are only sparingly soluble, because small

  \mathrm{Li^+}

  packs tightly with small or highly charged anions and gives a high lattice energy.
• Its halides have covalent character, so LiCl dissolves in organic solvents such as ethanol, and organolithium compounds are especially stable.
• It forms stable complexes such as

  \mathrm{[Li(NH_3)_4]^+}

  , which are rare for the heavier alkali metals.
• It does not form a solid hydrogencarbonate, and lithium hydroxide breaks down to

  \mathrm{Li_2O}

  and water on strong heating.

Thermal decomposition shows the effect of polarizing power most clearly. When a carbonate is heated, the cation pulls on the carbonate ion's oxygen and weakens a C–O bond until

\mathrm{CO_2}

breaks away. A tiny

\mathrm{Li^+}

does this well, and it also forms an oxide with a very high lattice energy, so lithium carbonate decomposes at temperatures where sodium and potassium carbonates are untouched.

[Li+].[Li+].[O-]C([O-])=O>>[Li+].[Li+].[O-2].O=C=O

Nitrates show the same split. Lithium nitrate decomposes to the oxide, nitrogen dioxide and oxygen, just as magnesium nitrate does, while sodium nitrate only loses oxygen to form the nitrite:

\begin{aligned} \mathrm{4LiNO_3} &\rightarrow \mathrm{2Li_2O + 4NO_2 + O_2} \\ \mathrm{2NaNO_3} &\rightarrow \mathrm{2NaNO_2 + O_2} \end{aligned}

How Beryllium Differs from the Other Alkaline Earth Metals

• It is very hard, hard enough to scratch glass, and has the highest melting point in group 2.
• A thin, tough oxide layer protects it, so it does not react with water even when red hot and resists attack by many acids.
• Its oxide and hydroxide are amphoteric, reacting with acids and also with strong alkalis.
• It is the only group 2 metal that dissolves in sodium hydroxide, releasing hydrogen.
• Its compounds are largely covalent:

  \mathrm{BeCl_2}

  is a poor conductor when molten and forms chains in which chlorine atoms bridge beryllium atoms.

\mathrm{Be + 2NaOH + 2H_2O \rightarrow Na_2[Be(OH)_4] + H_2}

The Diagonal Pairs Side by Side

Shared behaviour of the diagonal pairs

Property	Li and Mg	Be and Al
Reaction with N₂	Both form nitrides (Li₃N, Mg₃N₂)	Both form nitrides on strong heating
Oxide on burning	Normal oxide only (Li₂O, MgO)	Protective oxide film (BeO, Al₂O₃)
Carbonate on heating	Decomposes to oxide + CO₂	Carbonates unstable or absent
Oxide / hydroxide character	Basic	Amphoteric
Chlorides	Covalent character, soluble in ethanol	Covalent, bridged, fume in moist air
Reaction with alkali	None	Both release H₂

Charge density explains many trends across the periodic table, not just these anomalies. The same reasoning reappears when comparing how easily carbonates decompose down group 2 and why aluminium chloride is covalent while sodium chloride is ionic.

Tips for remembering

• Small ion, big pull: Li⁺ and Be²⁺ are tiny, so their charge is packed into a small space and they distort neighbouring anions, giving their compounds covalent character.
• Step right, step down: moving right raises charge density and moving down lowers it, so the diagonal neighbour (Li–Mg, Be–Al) ends up with a similar charge density and similar chemistry.
• Li acts like Mg: both make a nitride from nitrogen, form only a normal oxide in air, and have carbonates and nitrates that break down to the oxide on heating.