Main-Group Chemistry - Wyatt's Notes

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).

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.)

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).

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 $\text{IE}_2 \gg \text{IE}_1$.

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).

• 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).

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.

• 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.

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

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

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 $\text{CO}_3^{2-}$.

• 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.

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).

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.

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.

• 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.

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:

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

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).

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

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: $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$.

• 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).

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.

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: $\text{HOCl} < \text{HClO}_2 < \text{HClO}_3 < \text{HClO}_4$.

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

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.

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

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).

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

Down a group, oxides become more basic.

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

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.

flowchart TD
    A[Main Group Chemistry] --> B[Key Concepts]
    A --> C[Core Principles]
    A --> D[Practical Applications]
    B --> E[Fundamental definitions]
    C --> F[Design patterns]
    D --> G[Real-world usage]

• 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.

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).

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.

Problem: Sodium reacts vigorously with cold water to produce NaOH and H2. Predict and explain the trend in reactivity for the other alkali metals (Li, K, Rb, Cs) with water.

Solution: Reactivity with water increases down the group: Li < Na < K < Rb < Cs. The reaction is M(s) + H2O(l) -> MOH(aq) + 1/2 H2(g). The driving force is the ionization energy: as IE decreases down the group, it becomes easier to remove the valence electron and transfer it to water. Li reacts slowly (IE = 520 kJ/mol), Na reacts vigorously, K reacts violently (may ignite the hydrogen), Rb and Cs react explosively. The trend correlates directly with decreasing ionization energy and increasing atomic radius: larger atoms hold their outer electron less tightly, making them more reactive. The Standard Reduction Potential becomes more negative down the group (Li: -3.04 V, Cs: -2.92 V), confirming the trend.

Common mistake: Assuming that lithium is the most reactive alkali metal because it has the most negative reduction potential. While Li has the most negative E degree (-3.04 V), its high ionization energy and small atomic radius make it react more slowly with water than sodium. Reduction potentials account for sublimation energy, ionization energy, and hydration enthalpy, not just reactivity kinetics.

$\blacksquare$

Main-group chemistry encompasses the s-block and p-block elements, which together make up about 80% of the periodic table. The key insight is that trends repeat in predictable patterns. Down a group, atoms get larger and hold their electrons more loosely, making them more reactive (for metals) or less reactive (for nonmetals). Across a period, increasing nuclear charge pulls electrons closer, making atoms smaller and more electronegative. The inert pair effect is a relativistic phenomenon where heavy p-block elements prefer lower oxidation states because the ns^2 electrons are stabilized by relativistic contraction. Carbon is unique because it can form four strong covalent bonds and stable chains, giving rise to the entire field of organic chemistry. Silicon, its neighbor, cannot form stable double bonds because its larger 3p orbitals have poor overlap with oxygen’s 2p orbitals, leading to the formation of network solids like SiO2 instead of discrete molecules. The diagonal relationships (Li-Mg, Be-Al, B-Si) arise because elements diagonally adjacent have similar charge-to-radius ratios, producing similar polarizing power and chemical behavior.

• Thermodynamics
• Calculus