Structure and General Properties of the Nonmetals

Core Concepts

Which elements, and where

Nonmetals sit to the right of the metalloid staircase, plus hydrogen: C, N, O, the halogens, the noble gases, and (in group 15–16) P, S, Se. Their valence configurations run ns²np² through ns²np⁶, with H at 1s¹ and He at 1s². Nonmetallic character — the tendency to gain or share electrons — increases upward and to the right; fluorine is the most electronegative element of all.

Structure: small molecules and giant networks

Nonmetals bond covalently, and the number of valence electrons controls the structure:

• Seven diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂ — two atoms joined by single, double, or triple bonds (H–H, O=O, N≡N). The mnemonic "HOFBrINCl" (pronounced hof-brinkle) lists them.
• Polyatomic molecules: white phosphorus is P₄ (four atoms in a tetrahedron); sulfur forms S₈ rings.
• Monatomic gases: the noble gases exist as single atoms because their octets are already full.
• Network covalent solids: carbon (diamond, graphite) and silicon dioxide (quartz) bond in extended lattices — hard, high-melting, and the exceptions to "nonmetals are gases."

Because small molecules are held together only by weak intermolecular forces, most nonmetals melt and boil at low temperatures; network solids are among the hardest substances known.

Allotropes — different structural forms of the same element — are a nonmetal specialty: O₂ vs. ozone, diamond vs. graphite, white vs. red vs. black phosphorus, rhombic vs. monoclinic sulfur. Allotropes differ in properties even though the element is identical.

Bonding: anions and oxidation states

Nonmetals gain electrons from metals (forming anions like Cl⁻, O²⁻, N³⁻) or share them with each other. Because oxidation state is assigned as if electrons were transferred, a nonmetal can show negative states (bonded to less electronegative partners) or positive states (bonded to more electronegative partners such as O or F):

• Nitrogen: −3 in NH₃, +5 in HNO₃.
• Sulfur: −2 in H₂S, +6 in H₂SO₄.
• Chlorine: −1 in HCl, +7 in HClO₄.

The same element that grabs electrons from sodium hands them over to oxygen — a reversal that confuses students until they remember oxidation states are bookkeeping conventions.

Chemical behavior: acidic oxides and oxidizing power

• Acidic oxides: nonmetal oxides react with water to form acids — CO₂ → H₂CO₃, SO₂/SO₃ → H₂SO₃/H₂SO₄, NO₂ → HNO₃, P₄O₁₀ → H₃PO₄, Cl₂O₇ → HClO₄. This is the defining contrast with basic metal oxides and the origin of acid rain.
• Oxidizing strength: halogens are powerful oxidants in the order F₂ > Cl₂ > Br₂ > I₂, matching electronegativity; a stronger halogen displaces a weaker one from its salts: Cl₂ + 2Br^- ⟶ 2Cl^- + Br₂.
• Inertness of N₂: the N≡N triple bond (~945 kJ/mol) is so strong that nitrogen gas reacts only under extreme conditions (lightning, Haber process) — a blessing, since it keeps our atmosphere stable.
• Paramagnetic O₂: oxygen's two unpaired electrons make it attracted into a magnetic field — the only common paramagnetic gas.
• Noble gases: full octets make them nearly inert, but heavy ones do react with fluorine and oxygen (XeF₂, XeF₄, XeO₃) — "inert gas" is a myth of degree, not kind.

Examples: nonmetals by the numbers

Example 1 — Gas densities: why chlorine sinks

At STP, one mole of any ideal gas occupies 22.4 L. Density is mass per volume, so for any gas d = M/22.4 L. Nitrogen (M = 28.02 g/mol):

d(N₂) = 28.02 g22.4 L = 1.25 g/L

Chlorine (M = 70.90 g/mol):

d(Cl₂) = 70.90 g22.4 L = 3.17 g/L

Chlorine is about 2.5 times denser than air (~1.29 g/L), which is why a chlorine leak pools in low-lying areas — a real safety consideration whenever this toxic gas is used. (In the lab, chlorine is always handled in a fume hood.)

Example 2 — How much oxygen burns 5.00 g of sulfur?

Burning sulfur is the industrial route to sulfur dioxide, a precursor of sulfuric acid. The reaction: S₈ + 8O₂ ⟶ 8SO₂. Moles of sulfur first:

n(S₈) = 5.00 g8(32.07) g/mol = 5.00 g256.56 g/mol = 1.95 × 10^-2 mol

Now apply the mole ratio and convert to mass:

m(O₂) = 1.95 × 10^-2 mol S₈ × 8 mol O₂1 mol S₈ × 32.00 g O₂1 mol O₂ = 4.99 g

So 4.99 g of O₂ (and 10.0 g of SO₂ — mass is conserved: 5.00 + 4.99 ≈ 10.0 g).

Example 3 — Oxidation states of nonmetals: the bookkeeping rules

Oxygen is −2, hydrogen is +1 (with nonmetals), and the sum of oxidation states is zero for a neutral compound. Nitrogen in nitric acid, HNO₃:

1 + x + 3(-2) = 0   ⇒  x = +5

Sulfur in sulfuric acid, H₂SO₄:

2(1) + x + 4(-2) = 0   ⇒  x = +6

Chlorine in perchloric acid, HClO₄:

1 + x + 4(-2) = 0   ⇒  x = +7

The same nonmetals take negative states with less electronegative partners — N is −3 in NH₃ because x + 3(+1) = 0. Knowing these two directions lets you predict redox behavior without memorizing each compound.

Common Confusions