Occurrence & Electronic Configuration
Group 16 is termed the chalcogens — "ore-forming" — a name pointing to the fact that most copper ores contain either oxygen or sulphur and frequently the heavier members too. Oxygen is the most abundant element on earth, forming about 46.6% by mass of the crust, with dry air containing roughly 20.9% oxygen by volume. Sulphur is far scarcer (0.03–0.1% of the crust) and occurs as sulphates such as gypsum $\ce{CaSO4.2H2O}$ and baryte $\ce{BaSO4}$, and as sulphides such as galena $\ce{PbS}$, zinc blende $\ce{ZnS}$ and copper pyrites $\ce{CuFeS2}$.
Selenium and tellurium accompany sulphide ores as metal selenides and tellurides. Polonium occurs as a decay product of thorium and uranium minerals and is radioactive, while livermorium (Lv, Z = 116) is a synthetic, exceedingly short-lived element. All six members carry six electrons in the valence shell, giving the defining general configuration $ns^2\,np^4$.
| Element | Z | Outer configuration | Character |
|---|---|---|---|
| Oxygen (O) | 8 | [He] 2s²2p⁴ | Non-metal |
| Sulphur (S) | 16 | [Ne] 3s²3p⁴ | Non-metal |
| Selenium (Se) | 34 | [Ar] 3d¹⁰4s²4p⁴ | Metalloid |
| Tellurium (Te) | 52 | [Kr] 4d¹⁰5s²5p⁴ | Metalloid |
| Polonium (Po) | 84 | [Xe] 4f¹⁴5d¹⁰6s²6p⁴ | Metal (radioactive) |
The progression from non-metal (O, S) through metalloid (Se, Te) to metal (Po) is the usual top-to-bottom rise in metallic character. All these elements exhibit allotropy — oxygen as $\ce{O2}$ and ozone $\ce{O3}$, sulphur in rhombic and monoclinic forms.
Atomic Trends — Radii, Ionisation Enthalpy, Electronegativity
Down the group an extra principal shell is added at every step, so atomic and ionic radii increase from O to Po. Oxygen, sitting in the second period with no inner buffer, is exceptionally small — a fact that drives nearly all of its anomalous chemistry. Correspondingly, ionisation enthalpy decreases down the group as the outer electrons lie progressively further from the nucleus.
A subtle but heavily examined point: Group 16 elements have lower first ionisation enthalpies than the Group 15 elements of the same period. Group 15 has the extra-stable half-filled $np^3$ arrangement; in Group 16 the fourth p electron is forced to pair, and the electron–electron repulsion in that doubly-occupied orbital makes it relatively easy to remove. Electronegativity also falls down the group — oxygen has the second-highest electronegativity of all elements, behind only fluorine.
Moving down the chalcogens, size grows while the periodic "tightness" trends — ionisation enthalpy and electronegativity — fall away.
Electron gain enthalpy is not most negative at oxygen
Students assume the smallest atom must have the most negative electron gain enthalpy. In fact sulphur, not oxygen, has the more negative electron gain enthalpy. The oxygen atom is so compact that an incoming electron meets strong inter-electronic repulsion, making its $\Delta_{eg}H$ less negative than sulphur's. From sulphur onward the value again becomes less negative down to polonium.
Order of most-negative electron gain enthalpy: S > Se > Te > O > Po.
Oxidation States & the Anomalous Behaviour of Oxygen
Group 16 elements display oxidation states of −2, +2, +4 and +6, with +4 and +6 the most common for the heavier members. Because oxygen is so electronegative, it is essentially restricted to the −2 state, the only exceptions being $\ce{OF2}$ (where O is +2) and $\ce{O2F2}$ (where O is +1) — bonded to the one element more electronegative than itself. Sulphur, selenium and tellurium typically show +4 in their oxides and +6 with fluorine.
Down the group the stability of the +6 state decreases and that of the +4 state increases, a consequence of the inert-pair effect that grows heavier toward polonium. The −2 state also weakens going down; polonium hardly shows it at all.
Why oxygen's covalency is capped at four
Oxygen lacks accessible d orbitals, so it cannot expand its octet; its covalency is limited to four and in practice rarely exceeds two. Sulphur and its heavier congeners have vacant d orbitals available at low energy cost, so their octets expand freely — explaining why sulphur reaches valencies of 2, 4 and 6 in $\ce{H2S}$, $\ce{SF4}$ and $\ce{SF6}$ respectively, while no analogue of $\ce{SF6}$ exists for oxygen.
Strong hydrogen bonding in $\ce{H2O}$ (absent in $\ce{H2S}$) is the textbook signature of oxygen's small size and high electronegativity.
Hydrides — From H2O to H2Te
Every Group 16 element forms a hydride of type $\ce{H2E}$ (E = O, S, Se, Te, Po). These four examinable hydrides — $\ce{H2O}$, $\ce{H2S}$, $\ce{H2Se}$, $\ce{H2Te}$ — show beautifully regular gradations that NEET tests almost every year through statement-pairing questions. Three trends matter, and they all stem from the weakening of the H–E bond down the group.
| Property | Trend down the group | Reason |
|---|---|---|
| Thermal stability | $\ce{H2O}$ > $\ce{H2S}$ > $\ce{H2Se}$ > $\ce{H2Te}$ (decreases) | H–E bond enthalpy falls as E gets larger |
| Acidic character | $\ce{H2O}$ < $\ce{H2S}$ < $\ce{H2Se}$ < $\ce{H2Te}$ (increases) | Weaker H–E bond releases H⁺ more readily |
| Reducing power | $\ce{H2S}$ < $\ce{H2Se}$ < $\ce{H2Te}$ (increases) | Easier loss of H, water is non-reducing |
| Bond angle (H–E–H) | $\ce{H2O}$ (≈104.5°) > $\ce{H2S}$ > $\ce{H2Se}$ > $\ce{H2Te}$ (decreases) | Less s-character / weaker repulsion at larger central atom |
The aqueous acid dissociation constants make the acidity trend concrete: $K_a$ rises from about $1.8\times10^{-16}$ for water, through $1.3\times10^{-7}$ for $\ce{H2S}$ and $1.3\times10^{-4}$ for $\ce{H2Se}$, to $2.3\times10^{-3}$ for $\ce{H2Te}$. All the hydrides except water are reducing agents, the reducing character climbing from $\ce{H2S}$ to $\ce{H2Te}$.
Water (373 K) is anomalously high because of extensive intermolecular hydrogen bonding; from H₂S onward boiling point climbs with molar mass: H₂O > H₂Te > H₂Se > H₂S.
Boiling-point order is NOT a simple molar-mass series
A 2022 NEET item declared the order "$\ce{H2O} < \ce{H2S} < \ce{H2Se} < \ce{H2Te}$ because boiling point rises with molar mass" — and marked it incorrect. The true order is $\ce{H2O} > \ce{H2Te} > \ce{H2Se} > \ce{H2S}$. Hydrogen bonding lifts water far above the rest; only from $\ce{H2S}$ onward does the molar-mass trend apply.
Boiling point: $\ce{H2O}$ (373 K) ≫ $\ce{H2Te}$ > $\ce{H2Se}$ > $\ce{H2S}$.
The +4 and +6 oxides of sulphur lead straight into the oxoacids — see Sulphuric Acid for the contact process and the structure trends behind those S=O / S–OH bond counts.
Oxides & Their Classification
A binary compound of oxygen with another element is an oxide, and oxygen forms one or more oxides with nearly every element. Oxides are classified by their acid–base behaviour. Within Group 16 itself, the elements form $\ce{EO2}$ and $\ce{EO3}$ oxides (E = S, Se, Te, Po); both types are acidic, and the reducing tendency of the dioxides falls from $\ce{SO2}$ (reducing) to $\ce{TeO2}$ (oxidising).
| Type | Behaviour | Examples |
|---|---|---|
| Acidic | Combine with water to give an acid | $\ce{SO2}$, $\ce{CO2}$, $\ce{N2O5}$, $\ce{Cl2O7}$; high-oxidation-state metal oxides $\ce{Mn2O7}$, $\ce{CrO3}$, $\ce{V2O5}$ |
| Basic | Give a base with water | $\ce{Na2O}$, $\ce{CaO}$, $\ce{BaO}$ (generally metallic oxides) |
| Amphoteric | React with both acids and alkalies | $\ce{Al2O3}$, $\ce{ZnO}$ |
| Neutral | Neither acidic nor basic | $\ce{CO}$, $\ce{NO}$, $\ce{N2O}$ |
The defining reactions are clean to write in mhchem. An acidic oxide such as sulphur dioxide gives sulphurous acid:
$$\ce{SO2 + H2O -> H2SO3}$$
A basic oxide such as calcium oxide gives a hydroxide base:
$$\ce{CaO + H2O -> Ca(OH)2}$$
The amphoteric oxide alumina reacts both ways — with acid forming a salt, with alkali forming an aluminate:
$$\ce{Al2O3 + 6HCl + 9H2O -> 2[Al(H2O)6]^{3+} + 6Cl^-}$$
$$\ce{Al2O3 + 6NaOH + 3H2O -> 2Na3[Al(OH)6]}$$
Dioxygen — Preparation & Reactions
Dioxygen, $\ce{O2}$, is a colourless, odourless gas, sparingly soluble in water (about 3.08 cm³ per 100 cm³ at 293 K) — just enough to sustain aquatic life. It is uniquely paramagnetic despite having an even number of electrons, a fact explained by molecular orbital theory and the two unpaired electrons in its antibonding $\pi^*$ orbitals. In the laboratory it is generated by heating oxygen-rich salts or decomposing peroxides:
$$\ce{2KClO3 ->[MnO2][\Delta] 2KCl + 3O2}$$
$$\ce{2HgO ->[\Delta] 2Hg + O2}\qquad\ce{2Pb3O4 ->[\Delta] 6PbO + O2}$$
$$\ce{2H2O2 ->[MnO2] 2H2O + O2}$$
Industrially it is obtained by fractional distillation of liquefied air after removing $\ce{CO2}$ and water vapour, exploiting the difference in boiling points between dinitrogen and dioxygen. Dioxygen combines directly with most metals and non-metals (Au, Pt and the noble gases excepted). Because the $\ce{O=O}$ bond dissociation enthalpy is high (493.4 kJ mol⁻¹), reactions need initial heating but are then strongly exothermic and self-sustaining:
$$\ce{2Ca + O2 -> 2CaO}\qquad\ce{4Al + 3O2 -> 2Al2O3}$$
$$\ce{2ZnS + 3O2 -> 2ZnO + 2SO2}\qquad\ce{2SO2 + O2 ->[V2O5] 2SO3}$$
Ozone — Structure, Properties & Uses
Ozone, $\ce{O3}$, is the second allotrope of oxygen. At about 20 km altitude it forms from atmospheric oxygen under sunlight, building the ozone layer that screens the earth from harmful UV radiation. It is prepared in the laboratory by passing a slow, dry stream of oxygen through a silent electrical discharge (a Siemens ozoniser), which converts roughly 10% to ozone — "ozonised oxygen." The silent discharge is essential: a spark would decompose the product, and the process is endothermic.
$$\ce{3O2 -> 2O3}\qquad \Delta H^\circ (298\,\text{K}) = +142\ \text{kJ mol}^{-1}$$
Structure
Ozone is a bent (angular) molecule. The central oxygen uses $sp^2$ hybrid orbitals; both O–O bond lengths are identical at 128 pm — intermediate between a single (148 pm) and a double (121 pm) bond — and the bond angle is about 117°. This equal bond length is explained by treating ozone as a resonance hybrid of two equivalent structures, with the double bond delocalised between the two terminal atoms.
The two canonical forms differ only in which terminal bond is double; the real molecule averages them, giving identical 128 pm bonds and a 117° angle.
Properties & Oxidising Power
Pure ozone is a pale blue gas, a dark blue liquid (b.p. 161 K) and a violet-black solid. It is thermodynamically unstable with respect to oxygen: decomposition releases heat ($\Delta H$ negative) and increases entropy ($\Delta S$ positive), so $\Delta G$ for $\ce{O3 -> O2}$ is large and negative — high concentrations of ozone can be explosive. Because it readily liberates a nascent oxygen atom, ozone is a powerful oxidising agent:
$$\ce{O3 -> O2 + O}$$
$$\ce{PbS + 4O3 -> PbSO4 + 4O2}$$
$$\ce{2I^- + H2O + O3 -> 2OH^- + I2 + O2}$$
This last reaction underpins the quantitative estimation of ozone: ozone is passed into excess KI solution buffered at pH 9.2; the liberated iodine is then titrated against standard sodium thiosulphate. Ozone also reacts rapidly with nitric oxide, $\ce{NO + O3 -> NO2 + O2}$, the reaction at the heart of concerns about ozone-layer depletion by jet exhaust and freons.
In use, ozone serves as a germicide and disinfectant for sterilising water and purifying air, for bleaching oils, ivory, flour and starch, and as an oxidant in the manufacture of potassium permanganate.
Group 16 in one screen
- Configuration $ns^2np^4$; metallic character rises O → Po; all show allotropy.
- Down the group: radius ↑, ionisation enthalpy ↓, electronegativity ↓. Group 16 has lower IE than Group 15 (no half-filled stability); S has the most negative electron gain enthalpy, not O.
- Oxidation states −2, +2, +4, +6; oxygen capped at −2 (no d orbitals; +2 only in $\ce{OF2}$). +6 stability ↓, +4 stability ↑ down the group (inert-pair effect).
- Hydrides $\ce{H2E}$: thermal stability ↓, acidity ↑, reducing power ↑, bond angle ↓ from $\ce{H2O}$ to $\ce{H2Te}$. Boiling point $\ce{H2O} > \ce{H2Te} > \ce{H2Se} > \ce{H2S}$ (H-bonding lifts water).
- Ozone: bent, $sp^2$, both O–O = 128 pm, angle 117°, resonance hybrid; endothermic formation, powerful oxidiser via nascent O; estimated by KI–thiosulphate titration.