Main-Group Chemistry

1. Group 1: The Alkali Metals

1.1 Properties and Trends

Theorem 1 (Group 1 Trends):

• Melting and boiling points decrease down the group (weaker metallic bonding).
• Atomic and ionic radii increase.
• Ionization energy decreases: $\text{Li} > \text{Na} > \text{K} > \text{Rb} > \text{Cs}$.
• Electronegativity decreases.
• Reactivity increases (more vigorous reactions with water).

1.2 Key Compounds

Oxides:

• $\text{Li}_2\text{O}$ (normal oxide), $\text{Na}_2\text{O}_2$ (peroxide), $\text{KO}_2$ (superoxide).
• Stability of peroxide and superoxide increases down the group (larger cation stabilizes larger anion by lattice energy).

Hydrides: $\text{MH}$ (ionic, salt-like). Used as reducing agents.

Carbonates: $\text{M}_2\text{CO}_3$ — thermal stability increases down the group.

$\text{Li}_2\text{CO}_3 \xrightarrow{\Delta} \text{Li}_2\text{O} + \text{CO}_2$

(Na$_2$CO$_3$ is thermally stable.)

1.3 Lithium”s Anomalous Behavior

Lithium differs from other Group 1 elements due to its small size and high charge density:

• $\text{Li}_2\text{O}$ (not $\text{Li}_2\text{O}_2$).
• $\text{Li}_3\text{N}$ forms readily; other alkali metals do not.
• $\text{LiCl}$ is soluble in organic solvents (covalent character).
• Lithium resembles Mg (diagonal relationship).

2. Group 2: The Alkaline Earth Metals

2.1 Properties and Trends

Theorem 2 (Group 2 Trends):

• Harder, higher melting points than Group 1 (divalent metallic bonding).
• Be is amphoteric; Mg and heavier are basic.
• $\text{IE}_1$ and $\text{IE}_2$ both decrease down the group, but $\text{IE}_2 \gg \text{IE}_1$.

2.2 Key Compounds

Oxides: $\text{MO}$ (basic). BeO is amphoteric.

Hydroxides: $\text{M(OH)}_2$. Solubility increases down the group:

$\text{Be(OH)}_2 \ll \text{Mg(OH)}_2 < \text{Ca(OH)}_2 < \text{Sr(OH)}_2 < \text{Ba(OH)}_2$

Carbonates: $\text{MCO}_3$. Thermal stability increases down the group. All decompose on heating:

$\text{MCO}_3 \xrightarrow{\Delta} \text{MO} + \text{CO}_2$

Sulfates: Solubility decreases down the group ($\text{BaSO}_4$ is insoluble, used in X-ray imaging). $\text{MgSO}_4$ is soluble (Epsom salts).

2.3 Beryllium’s Special Properties

• Amphoteric oxide and hydroxide.
• Covalent bonding predominates (high charge density).
• Forms $[\text{Be(H}_2\text{O)}_4]^{2+}$ (tetrahedral, no octahedral complexes).
• Be resembles Al (diagonal relationship).

3. Group 13: Boron Group

3.1 Boron

Definition 1 (Boron): Metalloid with unique chemistry; forms covalent networks and electron-deficient compounds.

Boranes: $\text{B}_2\text{H}_6$ (diborane) features 3-center-2-electron bonds (banana bonds).

$\text{B}_2\text{H}_6 \text{ structure: } \text{Two BH}_2 \text{ units bridged by two H atoms}$

Boric acid: $\text{B(OH)}_3$ is a Lewis acid (not a Bronsted acid in the conventional sense):

$\text{B(OH)}_3 + \text{H}_2\text{O} \rightleftharpoons \text{B(OH)}_4^- + \text{H}^+$

Boron trihalides: $\text{BX}_3$ are strong Lewis acids, with strength $\text{BF}_3 < \text{BCl}_3 < \text{BBr}_3$. $\text{BF}_3$ is weaker than expected due to $p\pi$–$p\pi$ back-bonding from F lone pairs.

3.2 Aluminum

• Most abundant metal in Earth’s crust.
• Amphoteric: reacts with both acids and bases:

$2\text{Al} + 6\text{HCl} \to 2\text{AlCl}_3 + 3\text{H}_2$

$2\text{Al} + 2\text{NaOH} + 6\text{H}_2\text{O} \to 2\text{Na}[\text{Al(OH)}_4] + 3\text{H}_2$

• $\text{AlCl}_3$: Lewis acid catalyst (Friedel-Crafts); exists as $\text{Al}_2\text{Cl}_6$ dimer.

3.3 Group 13 Trend: Inert Pair Effect

Theorem 3 (Inert Pair Effect in Group 13): $+1$ oxidation state becomes more stable down the group:

$\text{Tl}^+ \text{ is more stable than } \text{Tl}^{3+}$

4. Group 14: Carbon Group

4.1 Carbon

Allotropes:

• Diamond: $sp^3$, tetrahedral network, hardest known material.
• Graphite: $sp^2$, layered sheets, excellent lubricant and conductor (within sheets).
• Fullerenes: C$_{60}$ (buckminsterfullerene), $sp^2$ with pentagonal rings.

Oxides:

• $\text{CO}_2$: Linear, nonpolar, greenhouses gas.
• $\text{CO}$: Toxic, strong ligand ($\sigma$ donor + $\pi$ acceptor).

Carbonates: $\text{H}_2\text{CO}_3$ (carbonic acid), bicarbonate $\text{HCO}_3^-$, carbonate $\text{CO}_3^{2-}$.

4.2 Silicon and Germanium

• Semiconductors (band gaps: Si 1.1 eV, Ge 0.67 eV).
• $\text{SiO}_2$: Network solid (silica), very different from CO$_2$.
• Silicates: Largest class of minerals; $\text{SiO}_4^{4-}$ tetrahedra share corners, edges, or faces.

4.3 Tin and Lead

Theorem 4: Inert pair effect pronounced:

• $\text{Sn}^{2+}$ (reducing agent, stannous) and $\text{Sn}^{4+}$ (stannic).
• $\text{Pb}^{2+}$ (more stable) and $\text{Pb}^{4+}$ (oxidizing agent).

Lead dioxide: $\text{PbO}_2$ is a strong oxidizing agent (used in lead-acid batteries).

5. Group 15: The Pnictogens

5.1 Nitrogen

Theorem 5 (Nitrogen Fixation):

$\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3 \quad \Delta H = -92 \text{ kJ/mol}$

Haber-Bosch process: High $T$, high $P$, Fe catalyst.

Oxides: N$_2$O (laughing gas), NO, N$_2$O$_3$, NO$_2$/N$_2$O$_4$, N$_2$O$_5$.

Oxides of nitrogen:

• NO: Radical (odd electron); biological signaling molecule.
• NO$_2$: Brown gas, odd electron.

Acids:

• $\text{HNO}_3$: Strong oxidizing acid; nitrates are soluble.
• $\text{HNO}_2$: Weak acid, unstable; nitrites.

5.2 Phosphorus

Allotropes: White P$_4$ (molecular, pyramidal), red P (polymeric), black P (layered).

Oxides: $\text{P}_4\text{O}_{10}$ (phosphorus pentoxide) is a powerful dehydrating agent.

Oxoacids:

• $\text{H}_3\text{PO}_4$ (phosphoric): Triprotic, pK$_a$ values: 2.15, 7.20, 12.35.
• $\text{H}_3\text{PO}_3$ (phosphorous): Diprotic (one H directly bonded to P).
• $\text{H}_3\text{PO}_2$ (hypophosphorous): Monoprotic.

5.3 Arsenic, Antimony, Bismuth

• Metalloids to metals down the group.
• $+5$ oxidation state becomes less stable; $+3$ dominates (inert pair effect).
• $\text{Bi}^{3+}$ is the common state; $\text{Bi}^{5+}$ is a strong oxidizer.

6. Group 16: The Chalcogens

6.1 Oxygen

Theorem 6: Most electronegative element after fluorine. Key compounds:

• $\text{H}_2\text{O}$: Anomalous (high boiling point, hydrogen bonding).
• $\text{H}_2\text{O}_2$: Peroxide; oxidizing and reducing agent.
• Ozone ($\text{O}_3$): Bent, resonance-stabilized, strong oxidant.

Theorem 7 (Chapman Cycle): Ozone formation and destruction in the stratosphere:

$\text{O}_2 \xrightarrow{h\nu} 2\text{O}$ $\text{O} + \text{O}_2 + \text{M} \to \text{O}_3 + \text{M}$ $\text{O}_3 \xrightarrow{h\nu} \text{O} + \text{O}_2$

6.2 Sulfur

Allotropes: S$_8$ (crown-shaped rings), polymeric sulfur at high $T$.

Oxides: $\text{SO}_2$ (bent, 119°), $\text{SO}_3$ (trigonal planar).

Oxoacids:

• $\text{H}_2\text{SO}_4$: Strong acid, strong dehydrating agent.
• $\text{H}_2\text{SO}_3$: Weak acid, sulfurous acid.

Sulfides: Metal sulfides have varying solubility; $\text{H}_2\text{S}$ is a weak acid (pK$_{a1}$ = 7.0).

6.3 Heavier Chalcogens

• Se and Te are semiconductors.
• Po is radioactive.
• Oxidation states range from $-2$ to $+6$; $+4$ and $+6$ dominate for S, Se, Te.

7. Group 17: The Halogens

7.1 Properties and Trends

Theorem 8 (Halogens Trends):

• Diatomic molecules (F$_2$, Cl$_2$, Br$_2$, I$_2$).
• State: gas (F$_2$, Cl$_2$) → liquid (Br$_2$) → solid (I$_2$).
• Electronegativity: F (3.98) > Cl (3.16) > Br (2.96) > I (2.66).
• Bond energy: $\text{F}_2 < \text{Cl}_2 > \text{Br}_2 > \text{I}_2$ (F$_2$ anomalously low due to lone pair repulsion).
• Reactivity decreases down the group: $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$.

7.2 Hydrogen Halides

• All are gases; HX bond strength decreases down the group.
• Acidity increases: HF (weak, pK$_a$ = 3.2) < HCl < HBr < HI (strong).
• HF is a weak acid despite high electronegativity (strong H–F bond and hydrogen bonding in solution).

7.3 Interhalogen Compounds

Definition 2 (Interhalogen): Compounds formed between two different halogens: XY, XY$_3$, XY$_5$, XY$_7$.

Examples: ClF, BrF$_3$, IF$_5$, IF$_7$.

The central atom is always the less electronegative halogen with the higher oxidation state.

7.4 Halogen Oxides and Oxoacids

Oxoacids of chlorine:

• $\text{HOCl}$ (hypochlorous): Weak acid, oxidizing agent (bleach).
• $\text{HClO}_2$ (chlorous): Weak acid.
• $\text{HClO}_3$ (chloric): Strong acid, strong oxidizer.
• $\text{HClO}_4$ (perchloric): Very strong acid, powerful oxidizer.

Acidity increases with oxidation state: $\text{HOCl} < \text{HClO}_2 < \text{HClO}_3 < \text{HClO}_4$.

8. Group 18: The Noble Gases

8.1 Properties

• All are monatomic gases.
• Very low boiling points (weak London dispersion forces).
• Full valence shells: extremely low reactivity.

8.2 Noble Gas Compounds

Theorem 9 (Noble Gas Reactivity): Only heavier noble gases form compounds:

• $\text{XeF}_2$, $\text{XeF}_4$, $\text{XeF}_6$: Fluorides of xenon.
• $\text{XeO}_3$, $\text{XeO}_4$: Oxides.
• $\text{KrF}_2$: Only krypton compound under extreme conditions.
• No true compounds of He, Ne, or Ar under normal conditions.

Xenon fluorides:

XeF$_4$ has 12 valence electrons (2 lone pairs on Xe); square planar by VSEPR.

9. Hypervalent Compounds

9.1 Definition and Examples

Definition 3 (Hypervalent): Molecules where the central atom has more than 8 valence electrons: $\text{PF}_5$, $\text{SF}_6$, $\text{ClF}_3$, $\text{XeF}_4$.

9.2 VSEPR for Hypervalent Compounds

9.3 3-Center-4-Electron Bonding Model

Theorem 10 (3c-4e Model): Hypervalent bonding is better described using 3-center-4-electron bonds rather than expanded octets. For example, in $\text{XeF}_2$:

$\text{F}^{–}\cdots\text{Xe}^+\cdots\text{F}^{–}$

Two electrons in the bonding orbital, two in a non-bonding orbital, and the Xe lone pairs remain in regular orbitals. This avoids invoking $d$-orbital participation (which is energetically unfavorable for period 2 elements).

10. p-Block Chemistry Patterns

10.1 Oxidation State Trends

10.2 Acid-Base Character of Oxides

Theorem 11: Across a period, oxides change from basic → amphoteric → acidic.

Down a group, oxides become more basic.

10.3 Allotropy

Many p-block elements exhibit allotropy: C (diamond/graphite/fullerene), P (white/red/black), S (S$_8$/polymeric), Se (gray/red/black).

Common Pitfalls

1. Confusing normal oxides, peroxides, and superoxides. Na forms Na$_2$O$_2$ (peroxide) and KO$_2$ (superoxide), but Li forms Li$_2$O (normal oxide). Fix: Larger cations stabilize larger anions; this is explained by lattice energy and ion size matching.
2. Wrong oxidation states for oxoacids. The oxidation state of the central atom in H$_3$PO$_3$ is +3 (not +5), because one H is directly bonded to P and is not ionizable. Fix: Count all electronegativity differences carefully.
3. Assuming all Group 14 compounds are like carbon. SiO$_2$ is a network solid, not gaseous like CO$_2$. Fix: Si forms $\sigma$ bonds but not $\pi$ bonds as readily; p$\pi$-p$\pi$ overlap is poor for larger atoms.
4. Ignoring the inert pair effect for heavy p-block elements. Tl$^+$ is more stable than Tl$^{3+}$; Pb$^{2+}$ is common. Fix: Apply the inert pair effect for all p-block elements below period 3.
5. Wrong VSEPR geometry for hypervalent molecules. XeF$_4$ is square planar, not octahedral (2 lone pairs occupy axial positions). Fix: Always count lone pairs when determining geometry.
6. Confusing acid strength trends. HF is a weak acid despite F being the most electronegative element. Fix: HF has a very strong H–F bond (high bond dissociation energy) and extensive hydrogen bonding.
7. Wrong fluoride bonding model. Using expanded octets (sp$^3$d$^2$ hybridization) is problematic for hypervalent compounds. Fix: The 3-center-4-electron model better describes hypervalent bonding.

Summary

• Group 1–2: s-block metals; reactivity increases down the group; oxides, hydrides, carbonates; diagonal relationships (Li/Mg, Be/Al).
• Group 13: Boron (electron-deficient, boranes, Lewis acid); Al (amphoteric); inert pair effect (Tl$^+$).
• Group 14: C allotropes (diamond, graphite); Si/Ge (semiconductors); Sn/Pb (inert pair effect).
• Group 15: N$_2$ fixation; phosphorus allotropes; oxoacids of P; inert pair effect (Bi$^{3+}$).
• Group 16: O$_3$, H$_2$O (anomalous), S allotropes; sulfur oxoacids.
• Group 17: F$_2$–I$_2$ reactivity trends; interhalogens; oxoacids of Cl.
• Group 18: Xe compounds (XeF$_2$, XeF$_4$, XeF$_6$); 3-center-4-electron bonding model.

Worked Examples

Example 1: Predicting Acid-Base Behaviour of Oxides

Problem: Classify the following oxides as acidic, basic, or amphoteric: Na2O, Al2O3, P4O10, SO3, MgO. Solution: Na2O: basic (Group 1 metal oxide, forms NaOH in water). Al2O3: amphoteric (Group 13, reacts with both acids and bases). P4O10: acidic (non-metal oxide of a high-oxidation-state element, forms H3PO4). SO3: acidic (non-metal oxide, forms H2SO4). MgO: basic (Group 2 metal oxide, forms Mg(OH)2). The trend across a period is from basic (left) to acidic (right).

Example 2: Silicon vs Carbon Chemistry

Problem: Explain why silicon does not form stable double bonds with oxygen (analogous to carbon dioxide), and why SiO2 forms a giant covalent lattice rather than discrete molecules. Solution: The Si=O pi bond is weaker than the C=O pi bond because silicon’s 3p orbitals have poor overlap with oxygen’s 2p orbitals (size mismatch and reduced p-p overlap). Instead, Si forms four single Si-O bonds, each of which is strong (partial d-p pi bonding provides additional stabilisation). This leads to a 3D network of SiO4 tetrahedra (quartz) rather than discrete SiO2 molecules. This is why silicon carbonyl analogues do not exist and why silica has a very high melting point.

Cross-References