Group 16 Elements: The Oxygen Family

Group 16 elements are oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po), all with the valence configuration $n{s}^{2}n{p}^4$; they are called chalcogens. Down the group they change from non-metals (O, S) through metalloids (Se, Te) to a radioactive metal (Po), and +4 becomes more stable than +6. This page covers the trends of group 16 elements and their key compounds: dioxygen, ozone, hydrogen peroxide, sulphur, ${\mathrm{H}}_2\mathrm{S}$, $\mathrm{S}{\mathrm{O}}_2$ and sulphuric acid. JEE Main and NEET test the trends and oxygen's unique behaviour; JEE Advanced also asks about the compounds.

1. The Oxygen Family at a Glance

Group 16 contains oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po); the synthetic element livermorium (Lv) also sits here. They are called chalcogens (Greek chalkos, ore, and genes, born), because a very large number of metal ores are oxides and sulphides. The first four are non-metals or metalloids; polonium is a radioactive metal.

Occurrence. Oxygen is the most abundant element on Earth: about 46.6% of the crust by mass and about 21% of dry air by volume. Sulphur makes up only 0.03-0.1% of the crust. It occurs as sulphates (gypsum $\mathrm{CaS}{\mathrm{O}}_4\cdot 2{\mathrm{H}}_2\mathrm{O}$, Epsom salt $\mathrm{MgS}{\mathrm{O}}_4\cdot 7{\mathrm{H}}_2\mathrm{O}$, baryte $\mathrm{BaS}{\mathrm{O}}_4$) and sulphides (galena PbS, zinc blende ZnS, copper pyrites $\mathrm{CuFe}{\mathrm{S}}_2$), as ${\mathrm{H}}_2\mathrm{S}$ in volcanoes, and in proteins, eggs, onion, garlic, mustard, hair and wool. Se and Te occur in sulphide ores; Po is a decay product in thorium and uranium minerals.

1.1 Profiles of the Five Elements

2. Trends in Physical Properties

Key data for the five elements:

• Atomic and ionic radii increase down the group as shells are added; oxygen is exceptionally small.
• Ionisation enthalpy decreases down the group. It is lower than for group 15 (O 1314 vs N 1402 kJ mol${}^{-1}$), because group 15 atoms have the extra-stable half-filled $n{p}^3$ subshell, while group 16 atoms lose a paired electron, which is pushed out by repulsion.
• Electron gain enthalpy: oxygen's is less negative than sulphur's. The added electron enters oxygen's small, crowded $2p$ subshell and is repelled. From sulphur onwards the value becomes less negative down the group.
• Electronegativity: oxygen is the second most electronegative element after fluorine; the value falls down the group (Se is a small exception).
• Metallic character increases: O and S are non-metals, Se and Te metalloids, Po a metal.
• Physical state: oxygen is a diatomic gas; the others are solids. Melting and boiling points rise down the group, with a big jump from O to S because oxygen is ${\mathrm{O}}_2$ while sulphur is ${\mathrm{S}}_8$. All the elements show allotropy.

3. Oxidation States and the Anomalous Behaviour of Oxygen

• Oxygen shows $-2$ in almost all its compounds, because it is second only to fluorine in electronegativity. Exceptions: +2 in $\mathrm{O}{\mathrm{F}}_2$, +1 in ${\mathrm{O}}_2{\mathrm{F}}_2$, $-1$ in peroxides (${\mathrm{H}}_2{\mathrm{O}}_2$, $\mathrm{N}{\mathrm{a}}_2{\mathrm{O}}_2$) and $-\frac{1}{2}$ in superoxides ($\mathrm{K}{\mathrm{O}}_2$).
• S, Se and Te show $-2$, +2, +4 and +6. The stability of $-2$ falls down the group. Their bonding in the +4 and +6 states is mainly covalent.
• Inert pair effect: down the group the +4 state becomes more stable than +6. $\mathrm{S}{\mathrm{O}}_2$ is a reducing agent, but $\mathrm{Te}{\mathrm{O}}_2$ is an oxidising agent; polonium shows mainly +2 and +4.

Why oxygen is different. Small size, high electronegativity and the absence of $d$ orbitals give oxygen its unusual behaviour:

• Hydrogen bonding: water is a liquid, while ${\mathrm{H}}_2\mathrm{S}$, ${\mathrm{H}}_2\mathrm{Se}$ and ${\mathrm{H}}_2\mathrm{Te}$ are gases (see Solved Example 4).
• Covalency: oxygen cannot exceed four (and rarely exceeds two), because it has no $d$ orbitals. Sulphur and the heavier elements use $d$ orbitals and reach six, as in $\mathrm{S}{\mathrm{F}}_6$.
• Multiple bonds: oxygen forms a $p\pi$-$p\pi$ double bond and exists as the diatomic, paramagnetic gas ${\mathrm{O}}_2$; sulphur prefers single bonds and forms ${\mathrm{S}}_8$ rings.

4. Chemical Reactivity of the Group

Oxygen and sulphur are both reactive, but reactivity decreases down the group; oxygen combines with almost all the elements.

4.1 Hydrides (${\mathrm{H}}_2\mathrm{E}$)

All the elements form volatile hydrides ${\mathrm{H}}_2\mathrm{O},{\mathrm{H}}_2\mathrm{S},{\mathrm{H}}_2\mathrm{Se},{\mathrm{H}}_2\mathrm{Te}$ and ${\mathrm{H}}_2\mathrm{Po}$. Water is made by burning hydrogen in oxygen; the others by the action of acids on metal sulphides, selenides and tellurides:

• Physical nature: water is a colourless, odourless liquid, essential to living things. The other hydrides are colourless, poisonous gases with very unpleasant smells. Volatility rises from ${\mathrm{H}}_2\mathrm{O}$ to ${\mathrm{H}}_2\mathrm{S}$ and then falls; water is the least volatile because of intermolecular hydrogen bonding.
• Acidity increases ${\mathrm{H}}_2\mathrm{O}<{\mathrm{H}}_2\mathrm{S}<{\mathrm{H}}_2\mathrm{Se}<{\mathrm{H}}_2\mathrm{Te}$. The E-H bond gets longer and weaker down the group, so the proton is released more easily. ${\mathrm{H}}_2\mathrm{S}$ is a weak diprotic acid.
• Reducing character: all except water are reducing agents, and the power increases from ${\mathrm{H}}_2\mathrm{S}$ to ${\mathrm{H}}_2\mathrm{Te}$, again because the E-H bond weakens.
• Thermal stability decreases from ${\mathrm{H}}_2\mathrm{O}$ to ${\mathrm{H}}_2\mathrm{Po}$ (bond enthalpy 463 down to 238 kJ mol${}^{-1}$).
• Bond angle decreases from $104.5^{\circ }$ to $90^{\circ }$. As the central atom becomes less electronegative, the bonding pairs sit farther from it and repel each other less, so the angle closes (Solved Example 3).

4.2 Halides

The elements form halides of the types $\mathrm{E}{\mathrm{X}}_6$, $\mathrm{E}{\mathrm{X}}_4$, $\mathrm{E}{\mathrm{X}}_2$ and the dimeric ${\mathrm{E}}_2{\mathrm{X}}_2$, mostly by direct combination:

• Stability of the halides falls in the order ${\mathrm{F}}^{-}>\mathrm{C}{\mathrm{l}}^{-}>\mathrm{B}{\mathrm{r}}^{-}>{\mathrm{I}}^{-}$. Only the fluorides form the hexahalides ($\mathrm{S}{\mathrm{F}}_6$, $\mathrm{Se}{\mathrm{F}}_6$, $\mathrm{Te}{\mathrm{F}}_6$): colourless, octahedral gases. $\mathrm{S}{\mathrm{F}}_6$ is exceptionally inert because its six F atoms shield the sulphur (steric protection); it is used as an insulating gas.
• Tetrafluorides ($\mathrm{S}{\mathrm{F}}_4$ gas, $\mathrm{Se}{\mathrm{F}}_4$ liquid, $\mathrm{Te}{\mathrm{F}}_4$ solid) are $s{p}^{3}d$ hybridised: a trigonal bipyramid with one equatorial lone pair, giving a see-saw shape.
• Dihalides ($s{p}^3$, bent) are known for all except selenium. The monohalides are dimeric (${\mathrm{S}}_2{\mathrm{F}}_2,{\mathrm{S}}_2\mathrm{C}{\mathrm{l}}_2,{\mathrm{S}}_2\mathrm{B}{\mathrm{r}}_2,\mathrm{S}{\mathrm{e}}_2\mathrm{C}{\mathrm{l}}_2,\mathrm{S}{\mathrm{e}}_2\mathrm{B}{\mathrm{r}}_2$) and disproportionate, for example $2S{\mathrm{e}}_2\mathrm{C}{\mathrm{l}}_2\to \mathrm{SeC}{\mathrm{l}}_4+3Se$.
• Oxygen forms only the fluorides $\mathrm{O}{\mathrm{F}}_2$ and ${\mathrm{O}}_2{\mathrm{F}}_2$; compounds such as $\mathrm{Cl}{\mathrm{O}}_2$, ${\mathrm{I}}_2{\mathrm{O}}_5$ and $\mathrm{C}{\mathrm{l}}_2{\mathrm{O}}_7$ are oxides of the halogens (Group 17). Sulphur forms ${\mathrm{S}}_2\mathrm{C}{\mathrm{l}}_2$, $\mathrm{S}{\mathrm{F}}_4$, $\mathrm{S}{\mathrm{F}}_6$ and others.

4.3 Oxides

• Monoxides (EO) are uncommon; SO exists only as a short-lived species.
• Dioxides $\mathrm{E}{\mathrm{O}}_2$ are formed by all the elements. $\mathrm{S}{\mathrm{O}}_2$ is a gas, $\mathrm{Se}{\mathrm{O}}_2$ a volatile solid, and $\mathrm{Te}{\mathrm{O}}_2$ and $\mathrm{Po}{\mathrm{O}}_2$ are non-volatile crystalline solids. Their reducing power falls down the group: $\mathrm{S}{\mathrm{O}}_2$ is a reducing agent, $\mathrm{Te}{\mathrm{O}}_2$ an oxidising agent.
• Trioxides $\mathrm{E}{\mathrm{O}}_3$ are formed by S, Se and Te; the best known is $\mathrm{S}{\mathrm{O}}_3$. In the solid it exists as a cyclic trimer $(\mathrm{S}{\mathrm{O}}_3{)}_3$ or as long chains cross-linked into sheets. $\mathrm{Se}{\mathrm{O}}_3$ is a cyclic tetramer and $\mathrm{Te}{\mathrm{O}}_3$ a solid with a network structure.
• Both $\mathrm{E}{\mathrm{O}}_2$ and $\mathrm{E}{\mathrm{O}}_3$ are acidic; they dissolve in water to give the oxoacids.

5. Dioxygen (${\mathrm{O}}_2$)

5.1 Preparation

In the laboratory, by heating potassium chlorate with a little manganese dioxide as catalyst, or by heating some metal oxides and peroxides:

On a large scale, oxygen is obtained by the fractional distillation of liquid air (after removing $\mathrm{C}{\mathrm{O}}_2$ and water vapour): nitrogen (b.p. 77 K) distils off first, leaving oxygen (b.p. 90 K). It is also a by-product of the electrolysis of water. When water is electrolysed to make heavy water, ordinary ${\mathrm{H}}_2\mathrm{O}$ decomposes faster, so the remaining water becomes richer in ${\mathrm{D}}_2\mathrm{O}$.

5.2 Properties

• Colourless, odourless gas; 3.08 cm${}^3$ dissolves in 100 cm${}^3$ of water at 293 K, enough to support aquatic life. It liquefies at 90 K and freezes at 55 K. It has three stable isotopes, ${}^{16}\text{O}$, ${}^{17}\text{O}$ and ${}^{18}\text{O}$.
• Paramagnetic, with two unpaired electrons, as molecular orbital theory predicts.
• Combines directly with nearly all metals and non-metals except some noble metals (Au, Pt) and some noble gases. The reactions often need heat to start but then continue by themselves.

Uses: respiration and combustion; oxyacetylene welding and cutting; steel manufacture; oxygen cylinders in hospitals, for mountaineers and at high altitude; liquid oxygen as the oxidiser in rocket fuels.

6. Simple Oxides

A binary compound of oxygen with another element is an oxide. Simple oxides are classified by how they behave with water, acids and bases:

• Acidic oxides combine with water to give an acid and react with bases. Most non-metal oxides are acidic ($\mathrm{S}{\mathrm{O}}_2,\mathrm{C}{\mathrm{O}}_2,{\mathrm{N}}_2{\mathrm{O}}_5,\mathrm{C}{\mathrm{l}}_2{\mathrm{O}}_7$), and so are the oxides of metals in high oxidation states ($\mathrm{M}{\mathrm{n}}_2{\mathrm{O}}_7,\mathrm{Cr}{\mathrm{O}}_3,{\mathrm{V}}_2{\mathrm{O}}_5$).
• Basic oxides give a base with water and react with acids; the oxides of active metals ($\mathrm{N}{\mathrm{a}}_2\mathrm{O},\mathrm{CaO},\mathrm{BaO}$) are basic.
• Amphoteric oxides react with both acids and bases, like $\mathrm{A}{\mathrm{l}}_2{\mathrm{O}}_3$ and ZnO.
• Neutral oxides are neither acidic nor basic: CO, NO and ${\mathrm{N}}_2\mathrm{O}$.
• Mixed oxides behave like a mixture of two oxides: $\mathrm{F}{\mathrm{e}}_3{\mathrm{O}}_4=\mathrm{FeO}\cdot \mathrm{F}{\mathrm{e}}_2{\mathrm{O}}_3$ gives both iron(II) and iron(III) salts.

7. Ozone (${\mathrm{O}}_3$)

Van Marum noticed the smell of ozone near electrical machines in 1785; Schonbein named it in 1840, and Soret established the formula ${\mathrm{O}}_3$ (1865), showing that ozone is an allotrope of oxygen. The ozone layer, mainly 15-35 km up in the stratosphere, absorbs harmful ultraviolet radiation from the Sun. Nitric oxide from supersonic jet exhaust and chlorine atoms from chlorofluorocarbons (CFCs) destroy it. Near the ground, ozone is decomposed by dust particles.

7.1 Preparation

Ozone is made in an ozoniser by passing a silent electric discharge through dry, cold oxygen. Some ${\mathrm{O}}_2$ molecules split into atoms, which add to other ${\mathrm{O}}_2$ molecules. The reaction is endothermic, so a silent (spark-free) discharge is used to avoid heating, which would decompose the ozone:

The product, ozonised oxygen, contains 5-10% ozone by volume. Pure ozone is obtained by passing ozonised oxygen through a spiral cooled in liquid air: ozone condenses (b.p. 161 K, $-112$ °C), and the liquid is fractionally distilled to remove dissolved oxygen.

Other ozone-forming reactions: (i) fluorine with water at low temperature gives oxygen containing ozone; (ii) ultraviolet light converts oxygen into ozone (this happens in the stratosphere); (iii) electrolysis of acidified water at a platinum anode with a high current density, where nascent oxygen combines with ${\mathrm{O}}_2$.

7.2 Physical Properties

• Pale blue gas with a strong, characteristic smell (like rotten fish when concentrated); it condenses to a dark blue liquid at 161 K and freezes to violet-black crystals at about 80 K.
• Heavier than air; slightly soluble in water but more soluble in turpentine oil, glacial acetic acid and $\mathrm{CC}{\mathrm{l}}_4$.
• Poisonous: small amounts cause headache and nausea. It is diamagnetic.

7.3 Chemical Properties

(a) Decomposition. Ozone is thermodynamically unstable. It decomposes at about 573 K, and at room temperature in the presence of $\mathrm{Mn}{\mathrm{O}}_2$, platinum black, silver or lead dioxide, which catalyse the decomposition:

(b) Oxidising nature. Ozone readily gives up an oxygen atom (${\mathrm{O}}_3\to {\mathrm{O}}_2+[\mathrm{O}]$), which makes it one of the strongest oxidising agents, second only to fluorine among common reagents:

In each of the following reactions one O atom of ozone oxidises the substance and ${\mathrm{O}}_2$ is released:

• Tailing of mercury: the mercury(I) oxide formed dissolves in the mercury, which then sticks to glass and loses its meniscus.
• Potassium ozonide, $\mathrm{K}{\mathrm{O}}_3$, is an orange solid containing the paramagnetic ${\mathrm{O}}_3^{-}$ ion.

Reactions in which all three O atoms are used (no ${\mathrm{O}}_2$ is released):

(c) Bleaching. Nascent oxygen from ozone oxidises coloured matter to colourless products. It bleaches oils, ivory, flour, starch, waxes and wood pulp.

(d) Ozonides and ozonolysis. Ozone adds across C=C bonds of unsaturated compounds to form ozonides. Warming an ozonide with water (usually with zinc dust, which destroys the ${\mathrm{H}}_2{\mathrm{O}}_2$ formed) splits it into carbonyl compounds; this is ozonolysis.

(e) Reaction with peroxides. Ozone and a peroxide reduce each other, releasing oxygen:

7.4 Structure, Tests and Uses

Ozone is bent: both O-O bonds are 128 pm, between the single bond (148 pm, as in ${\mathrm{H}}_2{\mathrm{O}}_2$) and the double bond (121 pm, as in ${\mathrm{O}}_2$), and the angle is about $117^{\circ }$. It is a resonance hybrid of two structures (Figure 5).

Tests: (i) strong characteristic smell; (ii) mercury loses its fluidity (tailing); (iii) an alcoholic solution of benzidine turns brown; (iv) an alcoholic solution of tetramethyl base turns violet; (v) starch-iodide paper turns blue.

Estimation: ozone liberates iodine from a buffered KI solution, and the iodine is titrated against standard sodium thiosulphate; this measures the ozone quantitatively (Solved Example 6).

Uses: germicide and disinfectant; sterilising drinking water; bleaching oils, ivory, flour and starch; oxidising agent in making $\mathrm{KMn}{\mathrm{O}}_4$; ozonolysis in organic chemistry.

8. Hydrogen Peroxide (${\mathrm{H}}_2{\mathrm{O}}_2$)

8.1 Preparation

(i) Laboratory method: cold, dilute sulphuric acid on sodium peroxide or hydrated barium peroxide:

Anhydrous $\mathrm{Ba}{\mathrm{O}}_2$ is not used, because the insoluble $\mathrm{BaS}{\mathrm{O}}_4$ coats it and stops the reaction. Phosphoric acid avoids this problem:

(ii) Electrolysis: cold sulphuric acid is oxidised electrolytically to peroxodisulphuric acid, which is hydrolysed and distilled:

(iii) Industrial method: auto-oxidation of 2-ethylanthraquinol. Air oxidises it to 2-ethylanthraquinone and ${\mathrm{H}}_2{\mathrm{O}}_2$; hydrogen with a Pd catalyst regenerates the anthraquinol. The net reaction is the catalytic union of ${\mathrm{H}}_2$ and ${\mathrm{O}}_2$, giving about 1% ${\mathrm{H}}_2{\mathrm{O}}_2$, which is concentrated by distillation under reduced pressure.

8.2 Physical Properties and Structure

• Pure ${\mathrm{H}}_2{\mathrm{O}}_2$ is a pale blue, viscous liquid (m.p. 272.4 K, b.p. about 423 K), miscible with water in all proportions. A 30% solution is sold as '100 volume' hydrogen peroxide.
• It decomposes slowly, faster in light and with catalysts (metal ions, metal powders, $\mathrm{Mn}{\mathrm{O}}_2$), so it is stored in wax-lined bottles in the dark, with urea as a stabiliser.
• Structure: non-planar, 'open book' (Figure 6).

8.3 Chemical Properties

(i) Decomposition (disproportionation, O goes from $-1$ to $-2$ and 0):

(ii) Oxidising agent in acid, ${\mathrm{H}}_2{\mathrm{O}}_2+2{\mathrm{H}}^{+}+2{\mathrm{e}}^{-}\to 2{\mathrm{H}}_2\mathrm{O}$:

and in alkali, ${\mathrm{H}}_2{\mathrm{O}}_2+2{\mathrm{e}}^{-}\to 2O{\mathrm{H}}^{-}$:

(iii) Reducing agent, ${\mathrm{H}}_2{\mathrm{O}}_2\to {\mathrm{O}}_2+2{\mathrm{H}}^{+}+2{\mathrm{e}}^{-}$, when it meets a stronger oxidant:

Restoring old paintings: the white lead in old oil paintings slowly turns into black PbS through ${\mathrm{H}}_2\mathrm{S}$ in the air. ${\mathrm{H}}_2{\mathrm{O}}_2$ oxidises it to white $\mathrm{PbS}{\mathrm{O}}_4$ and restores the colour.

8.4 Tests and Uses

• It liberates iodine from potassium iodide (faster with a trace of $\mathrm{FeS}{\mathrm{O}}_4$); the iodine turns starch blue.
• Acidified dichromate gives a deep blue colour of $\mathrm{Cr}{\mathrm{O}}_5$ (extractable into ether), which slowly fades as $\mathrm{Cr}{\mathrm{O}}_5$ decomposes to $\mathrm{C}{\mathrm{r}}^{3+}$ and ${\mathrm{O}}_2$: $\mathrm{C}{\mathrm{r}}_2{\mathrm{O}}_7^{2-}+4{\mathrm{H}}_2{\mathrm{O}}_2+2{\mathrm{H}}^{+}\to 2Cr{\mathrm{O}}_5+5{\mathrm{H}}_2\mathrm{O}$.
• A titanium(IV) salt in conc. ${\mathrm{H}}_2\mathrm{S}{\mathrm{O}}_4$ turns orange (pertitanic acid): $\mathrm{T}{\mathrm{i}}^{4+}+{\mathrm{H}}_2{\mathrm{O}}_2+2{\mathrm{H}}_2\mathrm{O}\to {\mathrm{H}}_2\mathrm{Ti}{\mathrm{O}}_4+4{\mathrm{H}}^{+}$.
• Uses: bleaching hair, cotton, wool and paper pulp; mild antiseptic (perhydrol); making sodium perborate and percarbonate for detergents; 'green' oxidant in industry and for treating effluents; restoring paintings.

9. Sulphur: Extraction and Allotropes

9.1 Extraction

(a) Frasch process, for underground deposits. Three concentric pipes are sunk into the sulphur-bearing rock. Superheated water (about 438 K, under pressure) pumped down the outer pipe melts the sulphur (m.p. about 392 K); hot compressed air sent down the innermost pipe forces a froth of molten sulphur, water and air up the middle pipe. The sulphur is 99.5-99.9% pure and practically free of As, Se and Te.

(b) From natural gas and refinery gases. ${\mathrm{H}}_2\mathrm{S}$ is separated by absorption in monoethanolamine and converted into sulphur. Part of it burns to $\mathrm{S}{\mathrm{O}}_2$, which is then reduced by more ${\mathrm{H}}_2\mathrm{S}$ over an $\mathrm{F}{\mathrm{e}}_2{\mathrm{O}}_3$ catalyst:

9.2 Allotropes

Sulphur shows allotropy to a remarkable degree. The three common forms are:

The temperature at which rhombic and monoclinic sulphur are in equilibrium, 369 K, is the transition temperature. Rings of 6 to 20 sulphur atoms have also been made; cyclo-${\mathrm{S}}_6$ has a chair form. At about 1000 K sulphur vapour is mostly ${\mathrm{S}}_2$, which is paramagnetic like ${\mathrm{O}}_2$.