1. The s-block and p-block: configurations and general trends
This chapter is for formative assessment only (not in the annual theory paper). s-block: groups 1 and 2, valence configuration and . p-block: groups 13 to 18, to . Properties vary regularly within a group, but the first member differs from the rest because of its small size, high ionisation enthalpy, high electronegativity and the absence of d orbitals (maximum covalency 4).
Down a group: one more shell is added, so atomic and ionic radii rise and first ionisation enthalpy generally falls. A cation is smaller than its atom (same nuclear charge, fewer electrons); an anion is larger. Group 1 values: atomic radius Li 152, Na 186, K 227, Rb 248, Cs 265 pm; IE₁ 520, 496, 419, 403, 376 kJ mol⁻¹.
Formulas
- Valence configurations of the s-block and p-block
Key terms
- s-block element
An element whose last electron enters an ns subshell (groups 1 and 2).
- p-block element
An element whose last electron enters an np subshell (groups 13 to 18).
- First-member anomaly
The first element of a group (Li, Be, B, C) differs from its heavier congeners because of small size, high ionisation enthalpy and no d orbitals.
2. Groups 1 and 2: ionisation enthalpy, flame colours, hydration and reducing character
Ions formed: group 1 gives M⁺ because IE₂ is very large (Na: IE₁ 496, IE₂ 4562 kJ mol⁻¹, the second electron coming from the neon core); group 2 gives M²⁺ (Mg: 737 and 1450). Group 2 has a higher IE₁ than group 1 in the same period.
Flame colours: Li crimson red, Na golden yellow, K violet, Rb red-violet, Cs blue; Ca brick red, Sr crimson, Ba apple green. Be and Mg give none because their tightly held electrons are not excited by an ordinary flame.
Hydration enthalpy of M⁺ (kJ mol⁻¹): Li −506, Na −406, K −330, Rb −310, Cs −276: smaller ions have higher charge density and hydrate more strongly.
Reducing character: all are strong reducing agents. E° (V): Li −3.05, Na −2.71, K −2.93; Be −1.97, Mg −2.36, Ca −2.84, Ba −2.92. Li has the most negative E° in group 1 because of its very high hydration enthalpy, even though its IE₁ is the highest.
Key terms
- Flame colour
Visible light emitted when an excited outer electron falls back to a lower level in a flame; characteristic of the metal.
- Hydration enthalpy
Enthalpy change when one mole of a gaseous ion dissolves in water to give the hydrated ion; more negative for smaller, more highly charged ions.
- Reducing agent
A species that loses electrons readily; the s-block metals are strong ones.
3. Anomalous lithium and beryllium; diagonal relationships
Lithium (small, high polarising power): forms Li₂O (not a peroxide) and Li₃N; (others give nitrites); Li₂CO₃ decomposes; no solid hydrogencarbonate; LiCl is deliquescent and soluble in ethanol. Li resembles Mg.
Beryllium: covalent compounds, covalency at most 4 (no d orbitals), amphoteric BeO and Be(OH)₂, passive towards concentrated nitric acid, carbide gives methane (). Be resembles Al. .
Diagonal relationship: across a period size falls and charge rises; down a group size rises; diagonally the effects roughly cancel. Li⁺ 76 pm and Mg²⁺ 72 pm; charge per radius of Be²⁺ (2/31 = 0.065) and Al³⁺ (3/53.5 = 0.056).
Formulas
- Charge-to-radius ratio of an ion
Key terms
- Diagonal relationship
Similarity in properties between an element of period 2 and the element one group to its right in period 3 (Li–Mg, Be–Al, B–Si).
- Polarising power
The ability of a small, highly charged cation to distort the electron cloud of an anion, giving covalent character.
- Amphoteric oxide
An oxide that reacts with both acids and bases (BeO, Al₂O₃).
4. Group 13: trends, inert pair effect and boron
Configuration ; oxidation states +3 and +1. Irregular trends: Ga (135 pm) is smaller than Al (143 pm); IE₁ goes 801, 577, 579, 558, 589 kJ mol⁻¹ (B to Tl), because 3d (and 4f) electrons shield poorly. Electronegativity 2.0, 1.5, 1.6, 1.7, 1.8.
Inert pair effect: the ns² electrons become reluctant to bond down the group, so +1 grows more stable: Tl⁺ is more stable than Tl³⁺ (Tl³⁺ is oxidising). Boron is a non-metal with covalent, electron-deficient compounds, maximum covalency 4, acidic B₂O₃ and a Lewis-acid BCl₃: . Aluminium is a metal with amphoteric Al₂O₃ and covalency up to 6.
Key terms
- Inert pair effect
The increasing reluctance of the ns² electron pair to take part in bonding in heavier p-block elements, so the oxidation state lower by two becomes more stable.
- Electron-deficient compound
A compound in which the central atom has fewer than eight electrons, such as BCl₃ (six).
- Lewis acid
An electron-pair acceptor.
5. Group 14: trends, catenation and inert pair effect
Configuration ; oxidation states +4 and +2. Covalent radius C 77, Si 118, Ge 122, Sn 140, Pb 146 pm; IE₁ 1086, 786, 761, 708, 715 kJ mol⁻¹; electronegativity 2.5, 1.8, 1.8, 1.8, 1.9. C and Si are non-metal and metalloid, Sn and Pb are metals.
Catenation: the tendency of atoms to link into chains and rings. Single-bond enthalpy C–C 348, Si–Si 297, Ge–Ge 260, Sn–Sn 240 kJ mol⁻¹: carbon is far the best, and its C–H (414) and C–O (358) bonds are also strong, so its chains are stable.
Inert pair effect: +4 is stable for C, Si, Ge; Sn(II) is reducing (); Pb(II) is stable and PbO₂ is oxidising ().
Key terms
- Catenation
The self-linking of atoms of an element to form chains and rings.
- Reducing agent (Sn²⁺)
Sn(II) is readily oxidised to Sn(IV).
- Oxidising agent (PbO₂)
Pb(IV) is readily reduced to the more stable Pb(II).
6. Anomalous carbon and the allotropes of carbon
Carbon vs silicon: C has no d orbitals (maximum covalency 4), forms pπ–pπ multiple bonds (C=C, C≡C, C=O) and shows unlimited catenation. Si reaches 6 () and forms no stable multiple bonds. CO₂ is a gas of small molecules; SiO₂ is a giant covalent solid. CCl₄ is not hydrolysed, whereas .
Allotropes of carbon (general):
- Diamond: sp³, 3-D network, hardest natural substance, insulator, density 3.51 g cm⁻³.
- Graphite: sp², flat hexagonal layers with weak forces between layers, soft, lubricant, conducts along the layers (delocalised electrons), density 2.26 g cm⁻³.
- Fullerene C₆₀: sp², hollow cage of 20 hexagons and 12 pentagons, a molecular solid.
| Allotrope | Hybridisation | Hardness | Conducts electricity? |
|---|---|---|---|
| Diamond | sp³ | Very hard | No |
| Graphite | sp² | Soft | Yes, along layers |
| Fullerene C₆₀ | sp² | Soft solid | Poorly |
Key terms
- Allotropes
Different structural forms of the same element in the same physical state.
- pπ–pπ bonding
Sideways overlap of p orbitals to give a π bond; C and N and O do this effectively, heavier elements do not.
- Fullerene
A molecular allotrope of carbon in the form of a closed cage, C₆₀ being the best known.
Exam tips
Common mistakes
- Explaining the most negative E° of lithium by its ionisation enthalpy (it is the hydration enthalpy).
- Saying that Be and Mg give flame colours, or that Li forms a peroxide.
- Forgetting that Ga is smaller than Al, or that Tl(I) is more stable than Tl(III).
- Claiming that silicon forms stable multiple bonds like carbon.
- Mixing up the hybridisation of diamond (sp³) and graphite (sp²).
How marks are usually awarded
- 1-mark items: a name, a colour, a trend or one reason.
- Anomaly questions (2–3 marks): small size and high charge density (1), no d orbitals or polarising power (1), a named example (1).
- Inert pair effect (2 marks): definition (1) + example with oxidation states (1).