Main Group Elements and Their Compounds

Position of the Main Group in the Periodic Table

Main group elements are those in the s- and p-blocks of the periodic table:

• s-block: Groups 1 and 2 (plus hydrogen and helium)
• p-block: Groups 13–18

They are called “main group” because:

• Their valence electrons occupy s and p orbitals only.
• Their properties change in a regular (periodic) way across periods and within groups.
• Many of the most common elements in the Earth’s crust, atmosphere, and living organisms belong to the main group (e.g. C, N, O, Si, P, S, Na, K, Ca).

Transition metals (d-block) and f-block elements are not part of the main group and are treated separately.

Key general trends that are especially visible in main group elements:

• Moving down a group:
• Atomic radius increases.
• Ionization energy generally decreases.
• Metallic character increases (elements become more metal-like).
• Moving across a period (left → right):
• Atomic radius decreases.
• Ionization energy and electronegativity generally increase.
• Metallic character decreases (change from metals to nonmetals).

These trends strongly influence which types of compounds main group elements form and what oxidation states they can adopt.

Groups of the Main Group and Their Characteristic Behavior

Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr)

Characteristic features:

• One valence electron ($ns^1$ configuration).
• Very reactive metals; reactivity increases down the group.
• Readily form +1 cations (e.g. Na⁺, K⁺).
• Strongly reducing agents (they easily lose their valence electron).

Typical compound types:

• Ionic halides: NaCl, KBr, etc.
• Oxides and peroxides:
• $$ 4\,\text{Li} + \text{O}_2 \rightarrow 2\,\text{Li}_2\text{O} $$
• $$ 2\,\text{Na} + \text{O}_2 \rightarrow \text{Na}_2\text{O}_2 $$
• Hydroxides: NaOH, KOH (strong bases in water).
• Carbonates and hydrogencarbonates: Na₂CO₃, NaHCO₃.

Hydrogen (also in Group 1) is special: it forms H⁺ and H⁻ in different contexts and behaves as a nonmetal; its chemistry is usually treated separately.

Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)

Characteristic features:

• Two valence electrons ($ns^2$ configuration).
• Less reactive than alkali metals but still reactive; reactivity increases down the group.
• Typically form +2 cations (e.g. Mg²⁺, Ca²⁺).

Typical compound types:

• Oxides: MgO, CaO (basic oxides).
• Hydroxides: Ca(OH)₂ (bases; generally less soluble than Group 1 hydroxides).
• Carbonates: CaCO₃ (often sparingly soluble, important in geology).
• Sulfates: CaSO₄, BaSO₄ (BaSO₄ almost insoluble, used as a contrast agent).

Beryllium is an outlier: it forms more covalent compounds and has notable toxicity.

Group 13: Boron Group (B, Al, Ga, In, Tl)

Characteristic features:

• Three valence electrons ($ns^2 np^1$).
• Boron: nonmetal; forms mainly covalent compounds.
• Aluminum: light metal, forms mostly ionic (but often partly covalent) compounds.
• Heavier congeners: increasingly metallic.

Typical compound types:

• Oxides: B₂O₃, Al₂O₃ (amphoteric, can react with acids and bases).
• Halides: BX₃ (e.g. BCl₃, electron-deficient and Lewis-acidic), AlCl₃ (also a Lewis acid).
• Hydrides and organometallics: BH₃ (exists as B₂H₆), AlH₃; organoboranes and organoaluminum compounds.

Important features:

• Electron deficiency in many B and Al compounds leads to unusual bonding (e.g. multicenter bonds in boranes).
• Al compounds often act as Lewis acids and catalysts.

Group 14: Carbon Group (C, Si, Ge, Sn, Pb)

Characteristic features:

• Four valence electrons ($ns^2 np^2$).
• Carbon and silicon: predominantly nonmetal/semimetal; strong tendency for covalent bonding.
• Tin and lead: more metallic; show +2 and +4 oxidation states.

Typical compound types:

• Oxides:
• CO and CO₂ (molecular, covalent).
• SiO₂ (giant covalent network solid).
• SnO/SnO₂, PbO/PbO₂ (ionic or largely ionic).
• Hydrides:
• CH₄ (methane) and extensive organic chemistry.
• SiH₄, GeH₄, SnH₄ (simpler, less stable hydrides).
• Halides: CCl₄, SiCl₄, SnCl₂, PbCl₂, etc.

Trends in oxidation states:

• C and Si: mainly +4.
• Down the group, the +2 state becomes more stable (inert pair effect), especially for Pb (Pb²⁺ more stable than Pb⁴⁺).

Group 15: Pnictogens (N, P, As, Sb, Bi)

Characteristic features:

• Five valence electrons ($ns^2 np^3$).
• Nitrogen: nonmetal, often forms multiple bonds (N≡N, C≡N).
• Phosphorus: nonmetal, forms extensive networks and polyatomic species (P₄, polyphosphates).
• As, Sb, Bi: progressively more metallic.

Common oxidation states:

• −3 (in many compounds with metals or more electropositive elements).
• +3 and +5 (especially P, As, Sb; Bi tends to +3).

Typical compound types:

• Hydrides: NH₃ (ammonia), PH₃ (phosphine), etc.
• Oxides: N₂O, NO, NO₂, N₂O₅, P₂O₃, P₂O₅, etc.
• Oxyacids and oxyanions: H₃PO₄ (phosphoric acid), NO₂⁻, NO₃⁻.

Trend: Going down the group, the ability to form strong multiple bonds (like N≡N) decreases; heavier elements favor single bonds and extended structures.

Group 16: Chalcogens (O, S, Se, Te, Po)

Characteristic features:

• Six valence electrons ($ns^2 np^4$).
• Oxygen: highly electronegative, supports many oxidation states in its compounds (−2, −1 in peroxides, +1, +2 in rare cases).
• Sulfur and heavier: form many allotropes and a wide variety of oxyacids and oxoanions.

Common oxidation states:

• −2 (in sulfides, oxides).
• +4 and +6 (in sulfites/sulfates, sulfurdioxide/sulfurdioxide derivatives, etc. for S; analogous states for Se, Te).

Typical compound types:

• Oxides and peroxides: H₂O, H₂O₂, SO₂, SO₃.
• Hydrides: H₂O, H₂S, H₂Se (decreasing bond strength and increasing acidity down the group).
• Oxyacids and oxyanions: H₂SO₃ (sulfurous acid), H₂SO₄ (sulfuric acid), SO₃²⁻, SO₄²⁻.

Trend: Down the group, the elements become less electronegative and more metallic; the stability of higher oxidation states (+6) decreases and lower ones become more favored.

Group 17: Halogens (F, Cl, Br, I, At)

Characteristic features:

• Seven valence electrons ($ns^2 np^5$).
• Very reactive nonmetals; strong oxidizing agents, especially F₂ and Cl₂.
• Exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).

Common oxidation states:

• −1 (as halide ions X⁻).
• Positive oxidation states (up to +7) in many oxo-species, especially for Cl, Br, I (e.g. ClO⁻, ClO₂⁻, ClO₃⁻, ClO₄⁻).

Typical compound types:

• Hydrogen halides: HF, HCl, HBr, HI (acids in water; HF is special in strength and bonding).
• Metal halides: NaCl, CaCl₂, FeCl₃, etc. (mostly ionic).
• Oxoacids and oxoanions: HClO, HClO₂, HClO₃, HClO₄.

Reactivity pattern:

• Oxidizing power decreases down the group (F₂ strongest, I₂ weakest).
• Acidity of hydrogen halides increases down the group (HF < HCl < HBr < HI).

Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

Characteristic features:

• Full valence shell ($ns^2 np^6$ except He: $1s^2$).
• Very low reactivity; gases under normal conditions.
• Used as inert atmospheres in many chemical processes.

Compounds:

• For Ne and Ar, only extremely few and unstable species.
• Kr and especially Xe form stable compounds with highly electronegative elements:
• Xenon fluorides: XeF₂, XeF₄, XeF₆.
• Xenon oxides and oxofluorides: XeO₃, XeO₄, XeOF₄.

These compounds demonstrate that even noble gases can participate in chemical bonding under suitable conditions.

Typical Oxidation States and Bonding Types in Main Group Compounds

Main group elements commonly exhibit oxidation states related to their valence electron count:

• Group 1: +1
• Group 2: +2
• Group 13: +3; heavier elements may also show +1 (inert pair effect).
• Group 14: +4 and +2 (down the group, +2 becomes more stable).
• Group 15: −3, +3, +5.
• Group 16: −2, +4, +6.
• Group 17: −1, +1, +3, +5, +7.
• Group 18: mostly 0; some positive states for heavier noble gases in compounds.

Bonding types:

• Ionic bonding:
• Typical for combinations of metals from Groups 1 and 2 with nonmetals from Groups 16 and 17 (e.g. NaCl, CaO).
• Covalent bonding:
• Dominant in compounds of nonmetals (C, N, O, F, etc.), such as CO₂, NH₃, CH₄.
• Polar covalent:
• Between elements with moderate electronegativity differences (e.g. H₂O, HCl).

Down a group, bonds involving heavier main group elements tend to be:

• Longer and weaker.
• More polarizable.
• More prone to forming networks or polymeric structures in the solid state.

Types of Compounds and Structural Motifs

Hydrides

Hydrides are compounds of elements with hydrogen. For the main group, three broad types appear:

1. Ionic (saline) hydrides:
• Formed by the most electropositive metals (e.g. NaH, CaH₂).
• Primarily contain H⁻ (hydride ion).
• React vigorously with water to release H₂.
2. Covalent hydrides:
• Formed by p-block nonmetals (e.g. CH₄, NH₃, H₂O, HF).
• Molecular compounds with specific geometries (e.g. tetrahedral CH₄, bent H₂O).
3. Metallic (interstitial) hydrides:
• More typical of transition metals; for main group, some borderline cases exist with heavy metals.

Hydride stability often decreases down a group (e.g. CH₄ is very stable, SnH₄ is much less so).

Oxides

Almost every main group element forms an oxide. General patterns:

• Basic oxides:
• Formed by electropositive metals (Groups 1 and 2).
• React with water to give bases (e.g. Na₂O → NaOH).
• Acidic oxides:
• Formed by nonmetals (e.g. CO₂, SO₃, P₄O₁₀).
• React with water to give acids (e.g. SO₃ + H₂O → H₂SO₄).
• Amphoteric oxides:
• Can react with both acids and bases (e.g. Al₂O₃, ZnO, SnO₂).
• Related hydroxides also show amphoteric behavior.

Across a period, the character of the oxides changes from basic (metal oxides) to amphoteric (metalloids) to acidic (nonmetal oxides).

Halides

Halides are compounds with halogen atoms (F, Cl, Br, I):

• Main group metals form mostly ionic halides:
• Example: NaCl, CaCl₂, AlCl₃ (though AlCl₃ shows covalent character in some phases).
• Nonmetals often form molecular covalent halides:
• Example: PCl₃, PCl₅, SiCl₄, CCl₄.

Trends:

• Fluorides are often more stable and more ionic.
• Down the halogen group, halides typically become less volatile and more polarizable.

Oxyacids and Oxyanions

Many nonmetallic main group elements form oxyacids and their conjugate base anions:

• Nitrogen: HNO₂ (nitrous acid), HNO₃ (nitric acid); NO₂⁻, NO₃⁻.
• Phosphorus: H₃PO₃, H₃PO₄; H₂PO₄⁻, HPO₄²⁻, PO₄³⁻.
• Sulfur: H₂SO₃, H₂SO₄; SO₃²⁻, SO₄²⁻.
• Halogens: HClO, HClO₂, HClO₃, HClO₄; ClO⁻, ClO₂⁻, ClO₃⁻, ClO₄⁻.

General tendencies:

• Higher oxidation state → stronger oxyacid (within a series of the same central atom).
• For halogens in oxyacids, electronegativity and number of oxygen atoms both influence acid strength.

Trends in Acidity and Basicity of Main Group Compounds

Several systematic patterns are especially visible among main group hydrides and oxides:

• Hydrides of Group 15–17 elements:
• Acidity increases down the group:
• NH₃ < PH₃ (weak acid behavior) and
• HF < HCl < HBr < HI (for halogen hydrides).
• This is due to decreasing bond strength and changes in bond polarity.
• Oxides and hydroxides:
• Across a period: basic oxide → amphoteric oxide → acidic oxide.
• Down a group: basic character generally increases for metals; acidity trends more complex for nonmetals.

These regularities are central for predicting the acid–base behavior of many inorganic compounds.

Allotropes of Main Group Elements

Some main group elements exist in multiple structural forms (allotropes), which have different physical and chemical properties despite being the same element:

• Carbon:
• Diamond: 3D network; very hard, insulating.
• Graphite: layered; soft, conducts electricity along layers.
• Other forms (e.g. fullerenes, nanotubes).
• Phosphorus:
• White phosphorus (P₄): molecular, very reactive, sensitive to light and heat.
• Red phosphorus: polymeric, less reactive.
• Black phosphorus: layered, more thermally stable.
• Sulfur:
• S₈ ring structures (orthorhombic and monoclinic forms).
• Various polymeric sulfur forms at high temperatures.

Allotropes often lead to different reactivities and uses, even though the element is the same.

Importance and Applications of Main Group Elements and Their Compounds

Main group elements and their compounds dominate many fundamental chemical processes and technologies:

• Energy and materials:
• Alkali and alkaline earth compounds in batteries, glass making, and construction (e.g. Na⁺ in electrolytes, CaCO₃ in cement).
• Silicon (Group 14) in semiconductor technology.
• Environment and geochemistry:
• Carbonates, silicates, phosphates, and sulfates shape rocks, soils, and waters.
• Nitrogen and sulfur oxides in air pollution and acid rain.
• Life and biochemistry:
• C, H, N, O, P, and S form the backbone of biomolecules.
• Na⁺, K⁺, Ca²⁺, Mg²⁺ are crucial for nerve impulses, muscle function, and structural roles (bones, chlorophyll).
• Industrial chemistry:
• Halogens and their compounds in disinfection, bleaching, and polymer production (e.g. PVC).
• Phosphates and sulfates in fertilizers and detergents.
• Noble gases as inert atmospheres in welding and in lighting.

These examples illustrate why the main group elements and their compounds form a central part of inorganic chemistry and why their periodic trends are so useful for prediction and understanding.