Q.Explain diagonal relationship.

Diagonal Relationship: Why Lithium Acts Like Magnesium

You know the periodic table as rows and columns. Elements in the same group share similar valence electron configurations, so they behave alike. Sodium and potassium are both reactive metals that form +1 ions. That much is straightforward.

But here is the surprise: lithium, the first element of Group 1, is far more similar to magnesium (Group 2) than it is to its own group-mate sodium. Drop a piece of lithium into water — it fizzes slowly. Drop sodium — it races across the surface, sometimes catching fire. Lithium's behaviour is closer to magnesium's: both react slowly with water, both form nitrides when heated in nitrogen, both give carbonates that decompose on heating.

This unexpected kinship is the diagonal relationship.

The Intuition: A Balancing Act

Imagine two properties that pull in opposite directions as you move across the periodic table.

• Going down a group: Atomic size increases, so the charge-to-size ratio (often called ionic potential, $\phi = \frac{\text{charge}}{\text{radius}}$) decreases. The element becomes more metallic, more reactive.
• Going right across a period: Atomic size decreases, and charge increases, so the charge-to-size ratio increases. The element becomes less metallic, more covalent in its bonding.

Now move diagonally: one step down and one step right. The increase in size (down) is roughly compensated by the increase in charge (right). The charge-to-size ratio stays nearly the same.

Lithium ($\text{Li}^+$, radius ~76 pm) and magnesium ($\text{Mg}^{2+}$, radius ~72 pm) have almost identical charge-to-size ratios. So does beryllium ($\text{Be}^{2+}$, ~31 pm) and aluminium ($\text{Al}^{3+}$, ~53 pm). The numbers are not exactly equal, but they are close enough to produce striking parallels.

The Precise Statement

The most important pairs are:

(You may also see C–P mentioned, but the similarity is weaker and rarely tested.)

Concrete Similarities: Li and Mg

Let's make this real with exam-relevant examples.

1. Reaction with nitrogen. Both lithium and magnesium burn in nitrogen to form nitrides.

Sodium and potassium do not form stable nitrides.

2. Carbonate decomposition. Lithium carbonate and magnesium carbonate both decompose on heating to the oxide and CO$_2$.

Other alkali metal carbonates (Na$_2$CO$_3$, K$_2$CO$_3$) are thermally stable — they do not decompose even at red heat.

3. Solubility of fluorides and phosphates. LiF and MgF$_2$ are both sparingly soluble in water. Li$_3$PO$_4$ and Mg$_3$(PO$_4$)$_2$ are also insoluble. The corresponding sodium and potassium salts are freely soluble.

4. Hydroxide behaviour. LiOH is a weaker base than NaOH or KOH. Mg(OH)$_2$ is a weak base too. Both are less soluble than the heavier alkali hydroxides.

5. Formation of organometallics. Both lithium and magnesium form alkyl and aryl compounds (organolithium reagents and Grignard reagents) that are used extensively in organic synthesis.

Be–Al Similarities

Beryllium and aluminium are another classic pair.

• Both are amphoteric — their oxides and hydroxides dissolve in both acids and bases.
• Both form covalent compounds (BeCl$_2$ is covalent, AlCl$_3$ is covalent in the anhydrous state).
• Both give carbides that produce methane on hydrolysis: Be$_2$C + 4H$_2$O → 2Be(OH)$_2$ + CH$_4$; Al$_4$C$_3$ + 12H$_2$O → 4Al(OH)$_3$ + 3CH$_4$.
• Both are rendered passive by concentrated nitric acid. …