`=>` Oxygen is the most abundant of all the elements on earth. Oxygen forms about `46.6%` by mass of earth’s crust. Dry air contains `20.946%` oxygen by volume.

`=>` The abundance of sulphur in the earth’s crust is only `0.03-0.1%`.

`=>` Combined sulphur exists primarily as sulphates such as gypsum `color{red}(CaSO_4.2H_2O)`, epsom salt `color{red}(MgSO_4.7H_2O)`, baryte `color{red}(BaSO_4)` and sulphides such as galena `color{red}(PbS)`, zinc blende `color{red}(ZnS)`, copper pyrites `color{red}(CuFeS_2)`.

`=>` Traces of sulphur occur as hydrogen sulphide in volcanoes.

`=>` Organic materials such as eggs, proteins, garlic, onion, mustard, hair and wool contain sulphur.

`=>` Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores.

`=>` Polonium occurs in nature as a decay product of thorium and uranium minerals.

`=>` The important atomic and physical properties of Group 16 along with electronic configuration are given in Table 7.6.

Some of the atomic, physical and chemical properties and their trends are discussed below.

The elements of Group 16 have six electrons in the outermost shell and have `color{red}(ns^2 np^4)` general electronic configuration.

`=>` Due to increase in the number of shells, atomic and ionic radii increase from top to bottom in the group.

`=>` The size of oxygen atom is, however, exceptionally small.

`=>` Ionisation enthalpy decreases down the group. It is due to increase in size.

`=>` The elements of this group have lower ionisation enthalpy values compared to those of Group 15 in the corresponding periods. This is due to the fact that Group 15 elements have extra stable half-filled `color{red}(p)`-orbitals electronic configurations.

`=>` Because of the compact nature of oxygen atom, it has less negative electron gain enthalpy than sulphur.

`=>` But from sulphur onwards the value again becomes less negative upto polonium.

`=>` Next to fluorine, oxygen has the highest electronegativity value amongst the elements.

`=>` Within the group, electronegativity decreases with an increase in atomic number.

`=>` This implies that the metallic character increases from oxygen to polonium.

Some of the physical properties of Group 16 elements are given in Table 7.6.

`=>` Oxygen and sulphur are non-metals, selenium and tellurium metalloids, whereas polonium is a metal.

`=>` Polonium is radioactive and is short lived (Half-life `13.8` days).

`=>` All these elements exhibit allotropy.

`=>` The melting and boiling points increase with an increase in atomic number down the group.

`=>` The large difference between the melting and boiling points of oxygen and sulphur may be explained on the basis of their atomicity; oxygen exists as diatomic molecule (`color{red}(O_2)`) whereas sulphur exists as polyatomic molecule (`color{red}(S_8)`).

Some of the chemical properties of Group 16 are given below :

The elements of Group 16 exhibit a number of oxidation states (Table 7.6).

`=>` The stability of `-2` oxidation state decreases down the group.

`=>` Polonium hardly shows `–2` oxidation state.

`=>` Since electronegativity of oxygen is very high, it shows only negative oxidation state as `–2` except in the case of `color{red}(OF_2)` where its oxidation state is + 2.

`=>` Other elements of the group exhibit `+ 2`, `+ 4`, `+ 6` oxidation states but `+ 4` and `+ 6` are more common.

`=>` Sulphur, selenium and tellurium usually show `+ 4` oxidation state in their compounds with oxygen and `+ 6` with fluorine.

`=>` The stability of `+ 6` oxidation state decreases down the group and stability of `+ 4` oxidation state increase (inert pair effect).

`=>` Bonding in `+4` and `+6` oxidation states are primarily covalent.

`=>` The anomalous behaviour of oxygen, like other members of `color{red}(p)`-block present in second period is due to its small size and high electronegativity.

`=>` One typical example of effects of small size and high electronegativity is the presence of strong hydrogen bonding in `color{red}(H_2O)` which is not found in `color{red}(H_2S)`.

`=>` The absence of `d` orbitals in oxygen limits its covalency to four and in practice, rarely exceeds two.

`=>` On the other hand, in case of other elements of the group, the valence shells can be expanded and covalence exceeds four.

(i) `color{green}("Reactivity with Hydrogen ")` All the elements of Group 16 form hydrides of the type `color{red}(H_2E (E = S, Se, Te, Po))`.

● Some properties of hydrides are given in Table 7.7.

● Their acidic character increases from `color{red}(H_2O)` to `color{red}(H_2Te)`. The increase in acidic character can be explained in terms of decrease in bond `color{red}(H–E)` dissociation enthalpy down the group.

● Owing to the decrease in bond `color{red}(H–E)` dissociation enthalpy down the group, the thermal stability of hydrides also decreases from `color{red}(H_2O)` to `color{red}(H_2Po)`.

● All the hydrides except water possess reducing property and this character increases from `color{red}(H_2S)` to `color{red}(H_2Te)`.

(ii) `color{green}("Reactivity with Oxygen ")` All these elements form oxides of the `color{red}(EO_2)` and `color{red}(EO_3)` types where `color{red}(E = S, Se, Te)` or `color{red}(Po)`.

● Ozone `color{red}(O_3)` and sulphur dioxide `color{red}(SO_2)` are gases while selenium dioxide `color{red}(SeO_2)` is solid.

● Reducing property of dioxide decreases from `color{red}(SO_2)` to `color{red}(TeO_2; SO_2)` is reducing while `color{red}(TeO_2)` is an oxidising agent.

● Besides `color{red}(EO_2)` type, sulphur, selenium and tellurium also form `color{red}(EO_3)` type oxides `color{red}(SO_3, SeO_3, TeO_3)`.

● Both types of oxides are acidic in nature.

(iii) `color{green}("Reactivity towards the Halogens ")` Elements of Group 16 form a large number of halides of the type, `color{red}(EX_6, EX_4)` and `color{red}(EX_2)` where `color{red}(E)` is an element of the group and `color{red}(X)` is a halogen.

● The stability of the halides decreases in the order `color{red}(F^(–) > Cl^(–) > Br^(–) > I^(–))`.

● Amongst hexahalides, hexafluorides are the only stable halides.

● All hexafluorides are gaseous in nature. They have octahedral structure.

● Sulphur hexafluoride, `color{red}(SF_6)` is exceptionally stable for steric reasons.

● Amongst tetrafluorides, `color{red}(SF_4)` is a gas, `color{red}(SeF_4)` a liquid and `color{red}(TeF_4)` a solid.

● These fluorides have `color{red}(sp^3d)` hybridisation and thus, have trigonal bipyramidal structures in which one of the equatorial positions is occupied by a lone pair of electrons. This geometry is also regarded as see-saw geometry.

● All elements except selenium form dichlorides and dibromides. These dihalides are formed by `color{red}(sp^3)` hybridisation and thus, have tetrahedral structure.

● The well known monohalides are dimeric in nature. Examples are `color{red}(S_2F_2, S_2Cl_2, S_2Br_2, Se_2Cl_2)` and `color{red}(Se_2Br_2)`. These dimeric halides undergo disproportionation as given below:

`color{red}(2Se_2Cl_2 → SeCl_4 + 3Se)`