The atomic-level trends in size and energy feed directly into how elements behave chemically — what valencies they show, how reactive they are, and whether their oxides are acidic or basic. This topic ties the physical periodicity to real chemistry.
Valence and oxidation states
The valence of a representative element usually equals the number of valence electrons (for groups 1, 2, 13, 14) or $8$ minus that number (for groups 15–17). Across period 3 the highest oxidation state rises with the group: Na (+1), Mg (+2), Al (+3), Si (+4), P (+5), S (+6), Cl (+7). Many p-block elements show variable oxidation states differing by 2 (e.g. Cl shows $-1, +1, +3, +5, +7$) because both $s$ and $p$ electrons can take part in bonding.
Anomalous properties of second-period elements
The first member of each group (Li, Be, B, C, N, O, F) differs from the rest because of (i) its small size, (ii) high electronegativity and charge density, and (iii) the absence of $d$ orbitals in the valence shell, which caps its covalency at 4. So nitrogen cannot expand beyond $\text{NCl}_3$ while phosphorus forms $\text{PCl}_5$; oxygen forms strong $p\pi$–$p\pi$ multiple bonds (as in $\text{O}_2$) while sulphur prefers chains and rings. This is linked to the diagonal relationship: Li resembles Mg, Be resembles Al, and B resembles Si because of similar charge-to-size ratios.
Periodic trends in chemical reactivity
Reactivity is highest at the two ends of a period and lowest in the middle. Metals (left) are most reactive when they lose electrons easily — reactivity increases down group 1 (Cs most reactive) as ionisation enthalpy falls. Non-metals (right) are most reactive when they gain electrons easily — reactivity increases up group 17 (F most reactive). The noble gases, with full shells, are almost inert.
Nature of oxides
The acid–base character of an oxide tracks the metallic character of the element. Across period 3 the oxides change from basic ($\text{Na}_2\text{O}$, $\text{MgO}$) through amphoteric ($\text{Al}_2\text{O}_3$) to acidic ($\text{P}_4\text{O}_{10}$, $\text{SO}_3$, $\text{Cl}_2\text{O}_7$). An amphoteric oxide reacts with both acids and bases, for example $\text{Al}_2\text{O}_3 + 6\text{HCl} \rightarrow 2\text{AlCl}_3 + 3\text{H}_2\text{O}$ and $\text{Al}_2\text{O}_3 + 2\text{NaOH} \rightarrow 2\text{NaAlO}_2 + \text{H}_2\text{O}$.
Metallic and non-metallic character
Metallic character (the tendency to lose electrons) decreases across a period and increases down a group, mirroring ionisation enthalpy. Non-metallic character shows the opposite trend. Between metals and non-metals lie the metalloids (B, Si, Ge, As, Sb, Te) along the diagonal staircase, sharing properties of both. So caesium is the most metallic stable element while fluorine is the most non-metallic.
Write the formulae of the highest oxides of Na, Al, Si, P, S and Cl in period 3 and state the trend in their acid-base nature.
Solution- The highest oxidation state equals the group's valence-electron count for these elements.
- Na (+1): $\text{Na}_2\text{O}$; Al (+3): $\text{Al}_2\text{O}_3$; Si (+4): $\text{SiO}_2$.
- P (+5): $\text{P}_4\text{O}_{10}$; S (+6): $\text{SO}_3$; Cl (+7): $\text{Cl}_2\text{O}_7$.
- Moving left to right, the metallic character falls, so oxides change from basic to acidic.
Answer: Na2O, Al2O3, SiO2, P4O10, SO3, Cl2O7; nature changes basic → amphoteric ($\text{Al}_2\text{O}_3$) → acidic.
Show with equations that aluminium oxide is amphoteric.
Solution- An amphoteric oxide reacts with both acids and bases.
- With an acid it behaves as a base: $\text{Al}_2\text{O}_3 + 6\text{HCl} \rightarrow 2\text{AlCl}_3 + 3\text{H}_2\text{O}$.
- With a base it behaves as an acid: $\text{Al}_2\text{O}_3 + 2\text{NaOH} \rightarrow 2\text{NaAlO}_2 + \text{H}_2\text{O}$.
- Since it reacts both ways, it is amphoteric.
Answer: $\text{Al}_2\text{O}_3$ reacts with HCl (as a base) and with NaOH (as an acid), confirming its amphoteric nature.
Why does nitrogen form NCl3 but not NCl5, whereas phosphorus forms both PCl3 and PCl5?
Solution- Nitrogen is a second-period element with valence shell $n=2$ (only $2s$ and $2p$ orbitals).
- It has no $d$ orbitals in its valence shell, so its maximum covalency is 4.
- Phosphorus ($n=3$) has empty $3d$ orbitals available for bonding.
- These $d$ orbitals let P expand its covalency to 5, forming $\text{PCl}_5$.
Answer: Lacking valence $d$ orbitals, N is limited to a covalency of 4 (NCl3), while P can use $3d$ orbitals to form PCl5.
State the diagonal relationship and give one pair of elements that exhibit it, with a reason.
Solution- The first element of a group often resembles the second element of the next group, placed diagonally.
- This happens because moving right increases charge density while moving down decreases it; the diagonal move roughly cancels, giving similar charge-to-size ratios.
- A classic pair is lithium and magnesium.
- Both form similar nitrides and burn in nitrogen, and their carbonates decompose on heating.
Answer: Li and Mg show a diagonal relationship because of their comparable charge-to-radius ratios (similarly Be–Al, B–Si).
Among Na, Mg, K and Ca, which is the most reactive metal and why?
Solution- Metallic reactivity depends on the ease of losing the valence electron, i.e. on a low ionisation enthalpy.
- Ionisation enthalpy decreases down a group and increases across a period.
- K is below Na in group 1 and group-1 metals lose their single electron most easily.
- So K has the lowest ionisation enthalpy among the four.
Answer: Potassium (K) is the most reactive, having the lowest ionisation enthalpy and hence losing its valence electron most readily.
How does metallic character vary across period 3 and down group 14? Give the reason.
Solution- Metallic character is the tendency to lose electrons, favoured by low ionisation enthalpy.
- Across period 3, $Z_{eff}$ rises and ionisation enthalpy increases, so metallic character decreases (Na metal → Cl non-metal).
- Down group 14, atomic size and shielding increase, lowering ionisation enthalpy.
- Hence metallic character increases down the group: C (non-metal) → Si, Ge (metalloids) → Sn, Pb (metals).
Answer: Metallic character decreases across period 3 and increases down group 14, tracking the change in ionisation enthalpy.