4_6 Group 16 (VIA) Elements: Comparing Oxygen and Sulphur

Group 16 (VIA) contains oxygen, sulphur, selenium, tellurium and polonium, together called the chalcogens, from the Greek for "ore-formers", because many metal ores are oxides or sulphides. Each atom has six valence electrons,

ns^2np^4

, and needs two more to reach a noble gas configuration. This is why −2 is the typical oxidation state in their compounds with metals. As in groups 14 and 15, the first element behaves quite differently from the rest.

Trends Down the Group

Group 16 elements: selected properties

Property	O	S	Se	Te	Po
Configuration	[He]2s²2p⁴	[Ne]3s²3p⁴	[Ar]3d¹⁰4s²4p⁴	[Kr]4d¹⁰5s²5p⁴	[Xe]4f¹⁴5d¹⁰6s²6p⁴
1st ionization energy (kJ/mol)	1314	1000	941	869	812
Electronegativity	3.4	2.6	2.6	2.1	2.0
Melting point (°C)	−218	115	221	450	254
Character	Non-metal	Non-metal	Non-metal	Metalloid	Radioactive metal

Going down, atoms grow, ionization energies and electronegativities fall, and metallic character rises, from non-metallic oxygen and sulphur to metallic polonium. Melting points jump from oxygen to sulphur, because oxygen is made of small

\mathrm{O_2}

molecules while sulphur is made of larger

\mathrm{S_8}

rings with stronger London forces. All members show allotropy: oxygen as

\mathrm{O_2}

and ozone, sulphur in rhombic, monoclinic and plastic forms, and selenium and tellurium in grey and red forms.

Oxygen's oxidation state is almost always −2, apart from peroxides (−1) and compounds with fluorine. The heavier elements also show +2, +4 and +6, as in

\mathrm{SCl_2}

,

\mathrm{SO_2}

and

\mathrm{SF_6}

. The older explanation was promotion of electrons into empty d orbitals. The modern view is that larger atoms can fit more neighbours around them, and that highly electronegative partners such as fluorine and oxygen can draw out all six valence electrons. Small oxygen can do neither.

Occurrence

Oxygen is the most abundant element in the Earth's crust, making up nearly half of it by mass, in silicates, oxides and carbonates. It makes up about 21% of the air by volume and 89% of water by mass. Sulphur occurs both free, in underground deposits and around volcanoes, and combined. Many metal ores are sulphides, such as galena (PbS), zinc blende (ZnS), cinnabar (HgS) and iron pyrites (

\mathrm{FeS_2}

). There are also sulphate minerals such as gypsum and barite. Most sulphur is now recovered from natural gas and crude oil. Sulphur is also part of proteins, through the amino acids cysteine and methionine, and gives onions, garlic and mustard their sharp smells.

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Oxygen and Sulphur: Similarities

• Both have the

  ns^2np^4

  outer configuration and form

  \mathrm{O^{2-}}

  and

  \mathrm{S^{2-}}

  ions with reactive metals.
• Both form covalent compounds with non-metals that have matching formulas, such as

  \mathrm{H_2O}

  and

  \mathrm{H_2S}

  , and

  \mathrm{CO_2}

  and

  \mathrm{CS_2}

  .
• Both are typical non-metals, both show allotropy, and both occur free and combined in nature.

Oxygen and Sulphur: Differences

Oxygen's small size lets its 2p orbitals overlap well side by side, forming a strong O=O double bond, so oxygen exists as small

\mathrm{O_2}

molecules. Sulphur's larger 3p orbitals overlap poorly sideways, so sulphur prefers two single bonds per atom and forms crown-shaped

\mathrm{S_8}

rings. These rings pack into a yellow solid. Oxygen also forms a second allotrope, ozone (

\mathrm{O_3}

), when an electric discharge or ultraviolet light passes through

\mathrm{O_2}

:

O=O.O=O.O=O>>[O-][O+]=O.[O-][O+]=O

Key differences between oxygen and sulphur

Property	Oxygen	Sulphur
State at room temperature	Gas (O₂)	Yellow solid (S₈ rings)
Allotropes	O₂ and O₃	Rhombic, monoclinic, plastic
Magnetism	Paramagnetic (two unpaired electrons)	Diamagnetic
Combustion	Supports burning	Burns to SO₂
Oxidation states	Mainly −2	−2, +2, +4, +6
Hydrogen bonding in hydride	Strong (H₂O is a liquid)	Negligible (H₂S is a gas)
Reaction with hot alkali	None	Forms sulphide and thiosulphate
Reaction with oxidizing acids	None	Oxidized to SO₂ or H₂SO₄

The hydride comparison is striking. Water boils at 100 °C because the highly electronegative oxygen forms strong hydrogen bonds. Hydrogen sulphide, with a heavier molecule, is a poisonous gas that boils at −60 °C, because sulphur is not electronegative enough to form them. Sulphur burns in air with a blue flame to form sulphur dioxide, the first step in making sulphuric acid.

S1SSSSSSS1.O=O.O=O.O=O.O=O.O=O.O=O.O=O.O=O>>O=S=O.O=S=O.O=S=O.O=S=O.O=S=O.O=S=O.O=S=O.O=S=O

Tips for remembering

• Oxygen double-bonds, sulphur chains: small oxygen forms strong O=O bonds and stays a diatomic gas, while larger sulphur forms single S–S bonds and builds S₈ rings.
• Oxygen stops at −2, sulphur goes to +6: oxygen is too electronegative to be oxidized by most elements, but sulphur can share all six valence electrons, as in SF₆ and H₂SO₄.
• O₂ is paramagnetic: it has two unpaired electrons, so liquid oxygen sticks between the poles of a magnet. Sulphur does not.