Where are the s- and p-blocks? (configuration)
Elements are placed in blocks by the orbital that receives the last electron.
- s-block: group 1 (alkali metals, ns¹) and group 2 (alkaline earth metals, ns²). All are reactive metals (hydrogen is a special case).
- p-block: groups 13 to 18, outer configuration ns² np¹ to ns² np⁶. It has metals (Al, Sn, Pb), metalloids (B, Si, Ge, As) and non-metals (C, N, O, F, noble gases).
Together they are the representative or main-group elements. Example: sodium (Z = 11) is [Ne] 3s¹; chlorine (Z = 17) is [Ne] 3s² 3p⁵.
Atomic and ionic radii
Down a group the radius increases: each new period adds a shell, and the inner electrons shield the outer ones from the nucleus.
Across a period the radius decreases: electrons are added to the same shell while the nuclear charge rises, so they are pulled in tighter.
A cation is smaller than its atom (Na⁺ < Na); an anion is larger than its atom (Cl⁻ > Cl). In group 13, Ga is about the same size as Al because the 3d electrons shield poorly.
Ionisation enthalpy and hydration enthalpy
Ionisation enthalpy
The energy needed to remove the outermost electron from a gaseous atom. It falls down a group (bigger atom, more shielding) and rises across a period. Group 1 metals have the lowest values, which is why they are such strong reducing agents. Small dips across a period: B < Be (2p electron easier to remove) and O < N (paired electron in 2p repels).
Hydration enthalpy
When an ion dissolves, water molecules surround it and energy is released. The smaller the ion and the higher its charge, the larger this energy: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺, and Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺. Group 2 ions (2+) hydrate much more than group 1 ions. This is why lithium salts are often hydrated (LiCl·2H₂O) and why lithium, despite a high ionisation enthalpy, is a very strong reducing agent in water.
Group and period trends at a glance
- Metallic character: increases down a group, decreases across a period.
- Electronegativity: decreases down a group, increases across (F is the highest).
- Oxides: s-block oxides are basic; across the p-block they turn amphoteric (Al₂O₃) and then acidic (SO₃, Cl₂O₇).
- Inert pair effect: in heavier p-block elements the ns² pair is reluctant to react, so lower oxidation states become more stable down a group (Tl⁺, Pb²⁺, Bi³⁺).
Why the first element is odd (anomalous behaviour)
Li, Be, B, C, N, O and F differ from the rest of their groups because they are very small, have high electronegativity and charge density, and have no d orbitals in the valence shell (maximum covalency 4).
- Li forms a normal oxide Li₂O and a nitride Li₃N, like Mg; its carbonate breaks up on heating.
- Be compounds are covalent; BeO is amphoteric, like Al₂O₃.
- C, N and O form strong pπ–pπ multiple bonds (C=C, N≡N, O=O); heavier members do not do this well.
- N cannot form NCl₅ but P forms PCl₅ (P has d orbitals).
Diagonal relationship: Li–Mg, Be–Al and B–Si look alike because moving right raises charge density and moving down lowers it, so the two effects nearly cancel.
Key formulas and definitions
- s-block outer configuration: ns¹ (group 1), ns² (group 2)
- p-block outer configuration: ns² np¹ to ns² np⁶ (groups 13–18)
- Group number for p-block = 10 + number of outer electrons
- Radius: increases down a group, decreases across a period
- Ionisation enthalpy: decreases down a group, increases across a period
- Hydration enthalpy: larger for smaller, more highly charged ions
- Diagonal pairs: Li–Mg, Be–Al, B–Si
Worked examples
1. Write the outer configuration and block of potassium (Z = 19).
K: [Ar] 4s¹. Last electron in an s orbital → s-block, group 1, period 4.
2. An element has outer configuration 3s² 3p³. Find its group, period and block.
Outer shell n = 3 → period 3. Electrons in p → p-block. Outer electrons = 5, group = 10 + 5 = 15. It is phosphorus.
3. Arrange Na, Mg, K in order of increasing atomic radius.
Mg is right of Na in the same period (smaller). K is below Na (larger). So Mg < Na < K.
4. Why is the first ionisation enthalpy of Mg (738 kJ/mol) higher than that of Al (578 kJ/mol)?
Al loses a 3p electron, which is higher in energy and shielded by the 3s pair. Mg must lose an electron from a filled 3s subshell, which is harder.
5. Which has larger hydration enthalpy, Na⁺ or Mg²⁺? Explain.
Mg²⁺. It is smaller (72 pm vs 102 pm) and has twice the charge, so it attracts water molecules much more strongly.
6. Give two reasons why lithium differs from the other alkali metals, and one way it resembles magnesium.
Reasons: very small atom and ion; high charge density (and high polarising power). Resemblance: both form nitrides by direct reaction with N₂ (Li₃N, Mg₃N₂), and both carbonates decompose on heating.
Common mistakes
- Thinking ionisation enthalpy rises smoothly across a period. It dips at group 13 (B < Be) and group 16 (O < N).
- Saying Li⁺ is the weakest reducer because its ionisation enthalpy is the highest. In water, its huge hydration enthalpy makes Li a very strong reducing agent.
- Putting helium in the s-block on the table but forgetting it is a noble gas placed in group 18 (it is 1s²).
- Mixing up "group number" in the p-block: it is 10 + outer electrons, not just the number of outer electrons.