Group – 16 elements Oxygen family
General Introduction :
→ Group-16 of periodic table has five elements. These elements are Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po). The elements of group-16 collectively known as oxygen family or chalcogen (meaning ore forming elements) because various ores of metals occur in the form of oxides and sulphides. In this family, first four members are non-metal while Polonium (Po) is radioactive element. Polonium was discovered by Marie Curie in 1998.
→ Occurence : Oxygen is the most abundant element of all the elements on earth. Oxygen forms about 46.6% by mass of earth’s crust. Dry air contains about 20.946% oxygen by volume.
→ On the other hand, the abundance of sulphur in the earth’s crust is about 0.03-0.1%. Sulphur mainly exist in the form of sulphates and sulphides. Some common examples are as follows :
- Gypsum – CaSO4 . 2H2O
- Epsom salt – MgSO4 . 7H2O
- Baryte – BaSO4
- Galena – PbS
- Zinc blendens – ZnS
- Copper pyrites – CuFeS2
- Cinnabar – HgS
- Iron pyrites – FeS2
- Celesite – SrSO4
→ Except sulphides and sulphates, traces of sulphur is also occur as hydrogen sulphide in volcanoes, in spring waters, coal gas, sewage gas, etc. Organic materials such as eggs, proteins, garlic, onion, mustard, hair and wool etc. also contains sulphur. Other elements of Group-16 as selenium, tellurium and polonium are found in the form of selenides and tellurides in sulphides ores.
→ Polonium occurs in nature as a decay product of thorium and uranium minerals. Its life span is very short. Oxygen gas with helium or carbon dioxide is used for artificial respiration. Liquid oxygen is also used as oxidiser for the rocket fuels.
→ Sulphur is used as disinfectant for destroying bacteria, fungi, insects etc. Sulphur is also used in the vulcanisation of rubber. It is also used in sulpha drugs for curing skin diseases. It is used to prepare SO2, H2SO4 gunpowder, fireworks etc.
Physical Properties Group – 16 elements
The main properties of group-16 are given in table 7.10 :
Table 7.10 : Some Physical Properties of Group-16 Elements
→ Electronic Configuration : The general electronic configuration of the elements of group-16 is nsnp. It indicates that last electron entres in p-sub shell. The electronic configuration of elements are given below :
Table 7.11 : Electronic Configuration
→ Atomic and lonic Radii : On moving down the group the atomic radius of elements increases because number of shells increases
O< S < Se < Te < Po (Atomic radii)
→ The atomic radii of the elements of group-16 are smaller than the corresponding elements of group-15 because on moving left to right in periodic table, nuclear charge increases due to greater attraction of electrons by the nucleus or due to increase in force of attraction between nucleus and valence electron. Atomic radii of > Elements of group-15 (Atomic radii) of elements of group-16
The size of anions also increases down the group due to increase in number of shells.
O2- < S2- < Se2- < Te2- < Po2- (Ionic radii)
→ Ionisation Energy: The ionisation enthalpies or ionisation energies of the elements of group-16 are high, thus the elements of this group do not loss electrons to form positive ions easily. As on moving down the groups the size of elements increases hence the amount of ionisation energy decreases.
O > S > Se > Te > Po (Ionisation enthalpy)
→ The first ionisation energy of the elements of group-16 is lower than the corresponding elements of group-15 because the elements of group-15 have half filled electronic configuration ie, its electronic configuration is ns np hence, they have higher value of ionisation energy.
→ While the second ionisation energy of the elements of group-16 is higher than the corresponding elements of group-15 because after losing one electron the configuration of elements of group-16 becomes half filled (ns2np3).
It can be understood with the values of ionisation energies of nitrogen and oxygen given below:
→ Electronegativity : The elements of group-16 are electronegative in nature, but on moving down the group electronegativity decreases due to increase in atomic size.
O > S > Se > Te > Po (Electronegativity)
Metallic Character On moving down the group metallic character increases due to increase in atomic size.
→ Density : The value of density increases on moving down the group because atomic weight increases comparatively more than atomic size.
O < S < Se < Te < Po (Density)
→ Melting and Boiling Points : Melting and boiling points increases down the group because on moving down the group molecular size increases. As a result, the magnitude to vander Waal’s forces increases and hence melting and boiling point also increases.
→ Catenation : It is the tendency of an atom to form bonds with indentical atoms. In group-16 only sulphur has strong tendency for catenation, Oxygen also shows this tendency but in a limited extent.
The order of catenation is as follows :
S >> 0 > Se > Te > Po (Catenation tendency)
→ As we know that sulphur shows maximum tendency of catenation, it is due to its high bond dissociation energy. Some examples of catenation are as follows :
→ Electron Affinity : The elements of group-16 have higher value of electron affinity. The value of electron affinity decreases from s to Po because on moving down the group atomic size increases and the distance between valence shell and nucleus increases. Therefore, additional electron which enters in valence shell experiences lesser force of attraction.
→ This force again decreases down the group so on moving down the group electron affinity decreases.
S > Se > Te > Po (Electron gain enthalpy)
→ The electron affinity of sulphur is greater than or more negative than oxygen. The low value of oxygen is due to its small size. As the size of oxygen is small, hence it has greater electron density which resists the addition of new electron.
→ The order of electron gain enthalpy is as follows:
O < S (Electron gain enthalpy)
→ Allotropy : All the elements of group-16 show allotropy. Allotrops of all the elements are as follows:
→ States of elements : At room temperature, oxygen is gas while other elements are solids. The size of oxygen is small hence it has a tendency to form multiple pπ – pπ bonds with other elements.
→ Due to formation of pπ – pπ multiple bonding, oxygen (O = O) molecule exist in gaseous state furthermore its molecules are held together by weak vander Waal’s forces hence it exists as gas at room temperature.
→ On the other hand, Sulphur (S), Selenium and Tellurium (Te) exists as solid at room temperature. They exist as staggered 8-atom rings with sp3 hybridisation on each atom. S8 molecule has puckered ring like or crown shaped structure.
→ Multiple Bonding : Except oxygen all the other elements of group-16 possess vacant d-orbitals in their valence shell. Hence these elements can form di-prt multiple bonding. Sulphur form comparatively stronger multiple bonding than other higher members of this group. For example,
CS2 (S = C = S) is more stable than
CSe2 (Se = C = Se).
→ Oxidation State : As we know that, the general electronic configuration of the elements of group-16 is nsnp hence, these elements have greater tendency to accept two electrons or share two electrons to achieve the noble gas configuration. These elements can show negative as well as positive oxidation state.
→ Negative Oxidation State : Oxygen is the highest electronegative element in group-16 (with ns2np4) hence it has maximum tendency to gain two electron to acquire noble gas configuration. Thus oxygen show-2 oxidation state in its compounds. Oxygen can also show.1 oxidation state in various peroxides.
Na2 O2 H2O2, BaO2 etc.
The tendency to show-2 oxidation state decreases down the group
→ Positive Oxidation State: Except oxygen other elements of group-16 show positive oxidation states. Oxygen show positive oxidation state only in O2E2 (oxidation state = +1), OF2 (oxidation state = +2).
→ Other elements of this group can show +2, +4 and +6 oxidation state in their compounds. It is due to the presence of empty d-orbitals in their shells. It can be explained in diagram shown below :
→ Hence it is clear from the above given diagram that the presence of unpaired electron decides the oxidation state. In ground state sulphur atom has two unpaired electron so it show +2 oxidation state. In first excited state one electron jumps from 3 porbitals to 3d-orbitals.
→ Now the number of unpaired electrons is 4 so it show +4 oxidation state. In second excited state one electron from 38 jumps into 3d. Now the number of unpaired electron is 6 hence it show +6 oxidation state. The tendency to show +2 oxidation state increases down the group due to inert pair effect.
Chemical Properties Group – 16 elements
Anomalous Behaviour of Oxygen :
→ The properties of oxygen are different from other elements. It is due to the following inherent characteristics
- It’s size is very small.
- It has very high ionisation energy.
- It has very high electronegativity.
- absence of d-orbitals in its valence shell.
→ The main points of difference between oxygen and remaining members of group-16 are given below:
- Oxygen is gas at room temperature while other are solids.
- Oxygen molecule is diatomic (O2) while other are polyatomic molecule ie, S8, Se8 with puckered ring like structure.
- Due to very high electronegativity of oxygen it is a typical non-metal. Sulphur is also non-metal but other members of this group have metallic character.
- Oxygen can form hydrogen bonding while others cannot. Due to the presence of hydrogen bonding, H2O exist as liquid while other hydrides are gaseous in nature.
- Oxygen exhibit 2 oxidation state only. In peroxides it show -1, in O2F2 (+1) and in OF2 it show +2 oxidation state. It does not show +4 and +6 oxidation state.
- Oxygen is the most abundant element in the earth’s crust.
- Oxygen is involved in multiple bonding (O = O, C = O etc.). But this tendency is absent in other elements.
- Compounds of oxygen are ionic as well as covalent while sulphur and other memebrs of this group generally forms covalent compounds.
- Oxygen use only porbitals in bonding while other can use d-orbitals also.
→ Trends in Chemical Reactivity Oxygen is the most reactive element of group-16. It can combine directly with almost all the elements i.e., metals as well as non-metals also. Its reactivity is slightly lesser than the halogens. The reactivity of the elements of group-16 decreases down the group
O > S > Se > Te > Po
Some important trends in the chemical reactivity of group-16 elements are discussed below:
→ Reactivity towards hydrogen (formation of hydride) All the elements of group-16 form hydrides of the type H2E (where E = 0, S, Se, Te, Po).
These hydrides are
→ Structure of Hydride: The central atom in hydride is sp3 hybridised, where two sp3 hybridised orbitals overlap with the orbitals of hydrogen and form two covalent bonds while other two sp3 hybridised orbital have lone pair of electrons. Hence, hydrides have angular or V-shaped structure as shown in figure 7.18.
→ Characteristics of Hydrides : The important characteristics of hydrides are given below:
→ Physical State : With an exception of water which is colourless, odourless liquid at room temperature, all other hydrides are pungent amelling gases.
→ Hydrogen Bonding : Oxygen due to its small size and high electronegativity forms hydrogen bonding. Except oxygen all the other elements of group-16 do not form hydrogen bonding due to low electronegativity and large size. Due to presence of hydrogen bonding H2O exist as liquid at room temperature as the molecules gets associated with each other.
→ Volatility: All the hydrogen are volatile. The volatility increases from H2O to H S then decreases down the group. Because on moving down the group atomic weight increases due to which vander Waal’s forces of attraction increases hence volatility decreases. The least volatility of H2O is due to the presence of hydrogen bonding.
H2O < H2S > H2Se > H2Te
→ Boiling and Melting Points : Boiling and melting points of hydrides increases down the group due to increase in vander Waals interaction. But the melting and boiling points of H2O is exceptionally high due to the formation of hydrogen bonding.
H2S < H2Se < H2Te < H2O
(Melting and boiling point)
→ Thermal Stability : Thermal stability of hydrides decreases down the group because atomic size increases hence the strength of H-E bond decreases so thermal stability decreases.
H2O > H2S > H2Se > H2Te (Thermal stability)
→ Reducing Character Except H,O all the other hydrides of the element of group-16 are reducing agent. Their reducing character increases down the group because thermal stability decreases.
H2O < H2S < H2Se < H2Te (Reducing character)
→ Covalent Character : Covalent character of these hydrides increases from H2O to H2 Te because as the electronegativity difference between element and hydrogen decreases, the covalent character increases, hence the water molecule has least covalent character but highest polarity. So we can say that covalent character is inverse of polarity.
H2O < H2S < H2Se < H2Te (Covalent character)
→ Bond Angle : On moving down the group the bond angle of hydride decreases because as soon as the atomic size of central atom increases the repulsive interaction between unpaired electron decreases hence the bond angle also decreases.
→ Acidic character : All the hydrides of this family are weak acidic in nature. The acidic · character increases down the group because the tendency to donate hydrogen increases.
H2O < H2S < H2Se < H2Te (Acidic character)
All the properties of hydrides of group-16 are summarised in table 7.12 :
Table 7.12: Properties of Hydrides of group-16 elements
→ Reactivity towards Oxygen (formation of Oxides) All the elements of group-16 reacts with oxygen and forms various types of oxides. These oxides are given in table 7.13.
Tebla 7.13: Oxides of elements of group-16
|Sulphur (S)||SO||SO2||SO3||S2O7, S2O, S2O3, SO4|
Oxides of the elements of group-16 are discussed briefly as below :
→ Mono-oxides : Except selenium all the elements of group-16 form monoxides. Structure of monoxide is linear. In structure
E = O (Here, E = S, Te, Po).
→ Dioxides : In oxides of elements of group-16, SO2 is a gas, SeO2 is a volatile solid and Teo, and PoO2 are crystalline solids. The acidic nature and stability of oxides decreases down the group. In dioxides, SO2 and SeO2 are acidic while TeO2 and PoO2 are amphoteric in nature. The acidic behaviour of dioxides of decreases down the group.
SO2 > SeO2 TeO2 (Acidic behaviour)
→ Structure : In dioxides the central atom is sp hybridised. It is angular in shape. Both the bonds in SO, is similar due to the presence of resonance. In dioxides, there are two types of bond i.e., Pπ – Pπ and pπ – dπ bond. Due to presence of lone pair of electrons the bond angle is slightly lesser than the normal bond angle (120°).
The structure of SO, is as follow :
SeO2 is crystalline solid. In solid state it exist as polymeric form but in gaseous state it exist as discrete monomeric form.
TeO2 and PoO2 are crystalline ionic solid. They are non-volatile.
Trioxides : The acidic behaviour of trioxides decreases down the group.
SO3 > SeO3 > TeO3 (Acidic behaviour)
Structure : In trioxides the central atom is sp2 – hybridised. Trioxides have planar triangular shape. SO3 molecule show resonance.
In solid state, SO3 exists as linear chain or a cyclic trimer.
Solid SeO3 also form cyclic structure. It exists as tetramer. However, in vapour state SeO3 has a monomeric structure like SO3.
The oxidation state and physical state of these oxides are given in table 7.13.
Table 7.13: Oxides of Sulphur
|Name||Formula||O.S. of S||Physical State|
|Sulphur suboxide||S2O||+1||Colourless gas|
|Sulphur mono oxide||SO||+2||Colourless gas|
|Sulphur sesquioxide||S2O3||+3||Bluish green crystalline solid|
|Sulphur dioxide||SO2||+4||Colourless gas|
|Sulphur trioxide||SO3||+6||Volatile liquid|
|Sulphur heptaoxide||S2O7||+7||Viscous liquid|
|Sulphur tetraoxide||SO4||+8||White solid|
Reactivity towards halogen (formation of halide)
The elements of group-16 form a number of halides. Some important halides are listed in table 7.14.
Table 7.14 : Halides of group-16
The properties of halide can be discussed as follows:
Monohalide : Some examples of monohalides are
S2F2, S2Cl2, O2F2, S2Br2, Se2Cl2, Se2 Br2 etc.
monohalides exist in dimeric form. They show disproportionation reaction. For example:
The structure of monohalide is similar to H2O2.
→ Dihalides : Except selenium all the other elements of group-16 form dihalides i.e. dibromide, dichloride. In dihalides the central atom is sp2 hybridised. Its structure is angular and geometry is tetrahedral. Due to the presence of two lone pair of electrons the bond angle is slightly less than 109.5°.
→ Tetrahalides: All the elements of group-16 form tetrahalides. Amongst tetrahalides, tetrafluorides are the most stable. The physical state of tetrahalides are different. Amongst tetrafluorides, SF4 is a gas, SeF4 liquid and TeF4 is solid. In tetrachlorides, SCl4 is unstable liquid while other are solids.
→ There are sp3 hybridisation in tetrahalides. It have trigonal bipyramidal structure in which one equatorial position is occupied by one lone pair of electrons. It has see-saw shape. Due to presence of lone pair-bond pair repulsion the bond angle decreases from 180° to 173°.
Hexahalides The stability of hexahalides decreases as mentioned below:
F > C > Br >I (Stability of halides)
→ Amongest hexahalides only hexafluorides are stable. All the hexahalides are gaseous in nature. The stability of hexafluorides again decreases as given below:
SF6 > SF6 > TeF6 (Stability)
→ SF6 is a colourless, odourless and non-toxic gas. It is thermally stable and chemically inert. The chemical inertness of SF, is due to the presence of six F-atoms. These six fluorine atoms protect the sulphur atom from attack by the reagents ie, sulphur atom is thermodynamically protected by six fluorine atoms.
All the hexafluorides are octahedral in shape in which sulphur atom undergoes sp3d2 hybridisation.