15.2 Representative Main-Group Chemistry: Alkali, Alkaline Earth & Halogens

15.2 Representative Main-Group Chemistry: Alkali, Alkaline Earth & Halogens

Quick Summary: Main-group descriptive chemistry emphasizes characteristic redox patterns, flame test spectra, and solubility trends. Group 1 alkali metals are soft, low-density reductants that form monoxides (Li2O), peroxides (Na2O2), or superoxides (KO2) based on cation-anion lattice matching. Group 2 alkaline earth metals exhibit opposing solubility trends for hydroxides (increasing down the column) versus sulfates (decreasing down the column), governed by the balance between lattice energy and cation hydration enthalpy. Halogens act as oxidizing agents whose strength decreases downward, undergoing temperature-dependent disproportionation in base. Noble gases form stable fluorides and oxides whose geometries follow VSEPR theory.

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1. Group 1: Alkali Metals (ns1)

The alkali metals (Li, Na, K, Rb, Cs) possess a single valence electron outside a noble gas core. They are soft, silvery metals with low densities (Li, Na, and K float on water) and low melting points that decrease down the column (Li: 180.5 °C; Cs: 28.5 °C).

Reducing Power and Water Reactivity

Possessing low first ionization energies and standard reduction potentials near -3 V (E°(Li+/Li) = -3.04 V due to Li+'s immense hydration enthalpy), alkali metals are formidable reducing agents. They react exothermically with water to liberate hydrogen gas and generate alkaline hydroxides:

2 M(s) + 2 H2O(l) → 2 M+(aq) + 2 OH-(aq) + H2(g)

Reactivity accelerates down the group: Li effervesces steadily; Na melts into a rolling sphere and often ignites; K bursts immediately into violet flame; Rb and Cs react explosively, shattering glass containers.

Combustion Products in Oxygen

Alkali metals form distinct oxides depending on cation-anion size matching and lattice energy:

• Lithium forms Monoxide (Li2O): Small Li+ stabilizes small oxide ion (O2-): 4 Li + O2 → 2 Li2O
• Sodium forms Peroxide (Na2O2): Intermediate Na+ stabilizes peroxide ([O-O]2-, formal oxidation state -1): 2 Na + O2 → Na2O2
• Potassium, Rubidium, Cesium form Superoxides (MO2): Large cations (K+, Rb+, Cs+) stabilize bulky superoxide (O2-, formal oxidation state -1/2): K + O2 → KO2. Potassium superoxide is used in rebreathers to absorb CO2 and generate O2:

4 KO2 + 2 CO2 → 2 K2CO3 + 3 O2

Flame Test Emissions

Thermal excitation of valence electrons produces characteristic visible emission lines: Li is crimson red (670.8 nm); Na is intense yellow (589.0, 589.6 nm, persistent D-lines); K is lilac / pale violet (766.5 nm, viewed through cobalt glass to filter trace sodium).

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2. Group 2: Alkaline Earth Metals (ns2)

Alkaline earth metals (Be, Mg, Ca, Sr, Ba) are harder, denser, and have higher melting points than Group 1 due to two valence electrons participating in metallic bonding. They display an exclusive +2 oxidation state. Flame tests yield diagnostic colors: Ca (brick-red), Sr (crimson/scarlet, highway flares), and Ba (apple-green). Magnesium burns with blinding white incandescence from hot particulate MgO.

Group 2 Compound Solubility Trends

Alkaline earth hydroxides and sulfates display opposing solubility trends down the group:

Solubility of Hydroxides:  Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2  (INCREASES DOWN)
Solubility of Sulfates:    MgSO4   > CaSO4   > SrSO4   > BaSO4    (DECREASES DOWN)

Thermodynamic Mechanism: Lattice vs Hydration Enthalpy

Dissolution enthalpy depends on lattice dissociation energy (ΔH°_lattice > 0) and hydration enthalpy (ΔH°_hydration < 0):

• Hydroxides (Small OH- Anion): Because OH- is small, lattice energy (proportional to 1/(r+ + r-)) drops sharply as cation radius r+ expands from Mg2+ to Ba2+. This rapid decrease outpaces the drop in cation hydration enthalpy. Consequently, dissolution becomes increasingly favorable down the group: Mg(OH)2 is insoluble (Ksp ≈ 5.6 × 10^-12, milk of magnesia), Ca(OH)2 is sparingly soluble (limewater), and Ba(OH)2 is a soluble strong base.
• Sulfates (Bulky SO4^2- Anion): Because SO4^2- is large, expanding r+ causes only a small fractional change in (r+ + r-), so lattice energy decreases slowly. Meanwhile, cation hydration enthalpy drops precipitously from Mg2+ (72 pm) to Ba2+ (135 pm). The loss of hydration stabilization dominates, making dissolution unfavorable: MgSO4 is soluble (Epsom salt), while BaSO4 is insoluble (Ksp ≈ 1.1 × 10^-10). In medicine, BaSO4 serves as a safe radiopaque gastrointestinal X-ray contrast agent because its insolubility prevents systemic barium poisoning.

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3. Group 17: Halogens (ns2 np5)

Halogens are reactive diatomic nonmetals with oxidizing power decreasing down the group: F2 > Cl2 > Br2 > I2. A higher halogen spontaneously oxidizes halide ions below it:

Cl2(aq) + 2 Br-(aq) → 2 Cl-(aq) + Br2(aq) (E°_cell = +0.29 V)

Br2(aq) + 2 I-(aq) → 2 Br-(aq) + I2(aq) (E°_cell = +0.53 V)

Base-Induced Disproportionation

Chlorine undergoes temperature-dependent disproportionation in aqueous hydroxide:

• Cold dilute base (0 to 20 °C): Cl2 + 2 OH- → Cl- + ClO- + H2O (household bleach)
• Hot concentrated base (70 to 100 °C): 3 Cl2 + 6 OH- → 5 Cl- + ClO3- + 3 H2O (chlorate)

Chlorine Oxoacids and Oxoanions Series

Chlorine forms four oxoacids: HClO (+1, hypochlorous), HClO2 (+3, chlorous), HClO3 (+5, chloric), and HClO4 (+7, perchloric). Acid strength increases with oxidation state (HClO < HClO2 < HClO3 < HClO4) due to inductive electron withdrawal by terminal oxygens and resonance delocalization of negative charge in the oxoanions.

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4. Group 18: Noble Gases (ns2 np6)

In 1962, Neil Bartlett observed that PtF6 oxidizes O2 to O2+[PtF6]-. Recognizing that xenon's first ionization energy (1170 kJ/mol) is comparable to that of O2 (about 1165 kJ/mol), Bartlett mixed Xe with PtF6 to synthesize the first noble gas compound, originally formulated as Xe+[PtF6]- (later shown to be a more complex mixture of xenon fluoroplatinates).

• Xenon Fluorides & Oxides (VSEPR):
  • XeF2: 5 electron pairs (2 bonding, 3 equatorial lone pairs) → Linear (180°).
  • XeF4: 6 electron pairs (4 bonding, 2 axial lone pairs) → Square planar (90°).
  • XeF6: 7 electron pairs (6 bonding, 1 lone pair) → Distorted octahedral.
  • XeO3: Formed by hydrolysis (XeF6 + 3 H2O → XeO3 + 6 HF); 4 electron domains (3 double bonds, 1 lone pair) → Trigonal pyramidal; shock-sensitive explosive solid.

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5. Main-Group Reference Tables

Main Group Reaction Profiles

Diagnostic Flame Test Identification

Chlorine Oxoacid Series