Group 1 and 2 elements — trends, anomalies, compounds, diagonal relationship

Question

Explain the diagonal relationship between lithium and magnesium. Why does lithium behave differently from the rest of Group 1? List three specific similarities between Li and Mg.

(CBSE Class 11 / NEET pattern)

Solution — Step by Step

Li (Group 1, Period 2) has similar properties to Mg (Group 2, Period 3) because their charge-to-size ratios (polarising power) are nearly equal.

Moving right across a period: charge increases, size decreases → polarising power increases. Moving down a group: size increases → polarising power decreases.

These two trends roughly cancel along the diagonal, giving Li and Mg similar polarising power.

flowchart LR
    A["Li\n(Group 1, Period 2)\nSmall size, +1 charge"] -->|"Diagonal\nrelationship"| B["Mg\n(Group 2, Period 3)\nLarger size, +2 charge"]
    C["Similar charge/size ratio\n→ Similar polarising power"] --> D["Similar chemical\nbehaviour"]
    A --> C
    B --> C

Both form nitrides directly with N₂: $6\text{Li} + \text{N}_2 \rightarrow 2\text{Li}_3\text{N}$ and $3\text{Mg} + \text{N}_2 \rightarrow \text{Mg}_3\text{N}_2$ . Other alkali metals do not form nitrides.
Both carbonates decompose on heating: $\text{Li}_2\text{CO}_3 \rightarrow \text{Li}_2\text{O} + \text{CO}_2$ and $\text{MgCO}_3 \rightarrow \text{MgO} + \text{CO}_2$ . Other Group 1 carbonates (Na₂CO₃, K₂CO₃) are stable to heat.
Both form covalent organometallic compounds: LiCH₃ and Mg(CH₃)₂ / Grignard reagents are both covalent. Other alkali metals form predominantly ionic organometallics.

Lithium has the smallest size and highest charge density of all alkali metals. This gives it:

Higher ionisation energy than other Group 1 elements
Greater tendency to form covalent bonds (high polarising power distorts electron clouds of anions)
LiF and Li₂O are much less soluble than NaF and Na₂O (small cation-small anion: high lattice energy)
Li does not form peroxides or superoxides (too small to stabilise large O₂⁻ or O₂²⁻ ions)

Property	Down Group 1	Down Group 2
Atomic radius	Increases	Increases
IE₁	Decreases	Decreases
Electronegativity	Decreases	Decreases
Melting point	Decreases	No clear trend
Reactivity with water	Increases	Increases
Oxide type	Li₂O → Na₂O₂ → KO₂	BeO → MgO → CaO → BaO₂

Why This Works

The diagonal relationship is a consequence of two competing periodic trends (across and down) approximately balancing each other. It is most prominent in Period 2 and 3 because the size differences are largest there. Li-Mg, Be-Al, and B-Si are the three classical diagonal pairs. Beyond these, the effect becomes negligible.

Alternative Method — Using Fajans’ Rules

Fajans’ rules predict when a bond has more covalent character: small cation, large anion, high charge. Li⁺ is small with high charge density, so it polarises anions more than Na⁺. This explains why LiCl is more covalent (soluble in organic solvents) while NaCl is strongly ionic.

For NEET MCQs, the two most tested facts are: (1) Li forms only the normal oxide ( $\text{Li}_2\text{O}$ ), Na forms peroxide ( $\text{Na}_2\text{O}_2$ ), K/Rb/Cs form superoxide ( $\text{KO}_2$ ). (2) Thermal stability of carbonates: Li₂CO₃ decomposes easily, Na₂CO₃ through Cs₂CO₃ are stable. The trend is driven by lattice energy — small Li⁺ strongly stabilises the small O²⁻ in Li₂O, making decomposition favourable.

Common Mistake

Students often write “Li forms Li₂O₂ (peroxide) like sodium.” This is wrong — lithium forms only the normal oxide $\text{Li}_2\text{O}$ , not peroxides or superoxides. The reason: Li⁺ is too small to stabilise the larger peroxide ( $\text{O}_2^{2-}$ ) or superoxide ( $\text{O}_2^-$ ) anions. Only larger cations like Na⁺, K⁺ can hold these bulky anions in a stable lattice.

Question

Solution — Step by Step

Why This Works

Alternative Method — Using Fajans’ Rules

Common Mistake

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