The P-Block Elements: Group 16 Elements

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Introduction 

The elements of the group 13 – 18 come under the p – block elements. In these elements the last electron enters in the outermost p – orbital. They have ns2np1-6 electronic configuration in valence shell, helium being an exception. These elements show the maximum oxidation state equal to the sum of electrons in the outermost shell or valence shell. Most of the elements of the p – block form covalent compounds although some elements form ionic compounds (such as halogens) and coordination compounds as well. p-block contains elements which are either metals, non – metals or metalloids. p-block elements include the group of halogens and inert gases. First member of each family of the p-block elements is given below in the table with their general electronic configuration and oxidation states. p-block has the most electronegative element which is fluorine. Elements of p-block generally form acidic oxides. Many elements such as C, Si, Ge, O, N etc. also show phenomenon of allotropy. Property of catenation is also shown by many elements.  


Group 

13

14

15

16

17

18

First Member of the Group

He

General Electronic Configuration 

ns2np1

ns2np2

ns2np3

ns2np4

ns2np5

ns2np6

Group Oxidation State 

+3

+4

+5

+6

+7

+8


We have covered the Boron Family (Group -13 elements), the Carbon Family (Group – 14 elements) and the Nitrogen Family (Group – 15 elements) in another article of p-block elements. In this article, we will cover the Oxygen Family or Group 16 Elements of p-block elements (Class XII, Chemistry). 


Group 16 is the fourth group of p-block elements. The first element of the group is oxygen, that’s why it is also known as the Oxygen Family. 

Group 16 Elements: The Oxygen Family 

Elements of the Group - 16

Atomic number 

Symbol 

Metal/nonmetal/metalloid

Color & State 

Electronic configuration 

Density g/cm3 at 298 K

Atomic 

and 

ionic radii 

Ionization enthalpy 

8

O

Non - metal 

Colorless gas 

[He] 2s2 2p4

1.32

Increases on moving from top to bottom in the group due to increase in the number of shells. 

Decreases on moving from top to bottom in the group due to gradual increase in size of elements.


16

S

Non- Metal 

Lemon colored crystalline solid  

[Ne] 3s2 3p4

2.06

34

Se

Non – metal 

Black, red and gray colored allotropes are found.

Solid 

[Ar] 3d10 4s2 4p4

4.19

52

Te

Metalloid

Silvery gray with luster. 

Solid 

[Kr] 4d10 5s2 5p4

6.25

84

Po

Metal

Silver colored.

Solid 

[Xe] 4f14 5d10 6s2 6p4

-


Elements of the Group 16 – Physical Properties 

Symbol 

Atomic number 

Atomic mass 

(g mol-1)

Melting 

point (K)

Boiling 

point (K)

Density 

Ionic 

radius 

Electro negativity 

O

8

16

55

Increases on moving from top to bottom (Exception – Po)

90

Increases on moving from top to bottom in the group(Exception – Po)

Increases on moving from top to bottom in the group 

Increases on moving from top to bottom in the group.

Decreases on moving from top to bottom in the group. 

Oxygen is the highest electronegative element after fluorine. 

S

16

32

393

718

Se

34

78.96

490

958

Te

52

127.60

725

1260

Po

84

210

520

1235

  • Oxygen has less negative electron gain enthalpy than other elements of the group due to its small size. 

  • All elements of the group 16 exhibit allotropy. 


Elements of the Group 16 – Chemical Properties

Oxidation Number 

Group -16 elements generally exhibit -2, +4 and +6 oxidation states.


As we move top to bottom in the group metallic character increases, so the stability of -2 oxidation state decreases. 

Oxygen exhibits -2 oxidation state. S, Se, Te show + 4 and + 6 oxidation states. Po shows mainly +4 oxidation state. 

Reactivity Towards Oxygen  

All elements of oxygen family forms two types of oxides – EO2, EO3

Both types of oxides are acidic in nature. 

Reducing property of dioxide decreases from S to Te.  

Reactivity with Hydrogen 

All elements of the oxygen family react with hydrogen and form H2E type hydrides where E = any element of group 16. 

Reducing character of the hydrides increases on moving down the group. Water is an exception. 

Acidic character of the hydrides of group 16 elements increases on moving down the group as the bond enthalpy of H-E bond decreases on moving down the group. 

Reaction Towards Halogens 

All elements of the oxygen family react with halogens. 

Group – 16 elements form following three types of halides – EX6, EX4 and EX2 where E = any element of group 16 and X = halogens 

The stability of the halides of group 16 decreases in the order F- > Cl- > Br- > I-.

All hexafluorides of the group – 16 elements are gaseous. 

All elements except oxygen form dichlorides and dibromides


Anomalous Properties of Oxygen 

  • Oxygen differs from other elements of the group – 16 due to its high electronegative character, small size and high ionization enthalpy. 

  • Hydride of oxygen means water molecules form hydrogen bonds due to high electronegative character of oxygen. 

  • d- orbitals are not found in the valence shell of the oxygen atom. It limits its covalency to four. While in case of other elements they have d – orbitals and their covalence exceed four. 


Compound 

Preparation 

Properties 

Uses 

Dioxygen 

🡪 Dioxygen is prepared by heating chlorates, nitrates and permanganates in presence of MnO2. 

Reaction of potassium chlorate –

2KClO3 ΔMnO2 ⟶ 2KCl + 3O2

🡪 By the thermal decomposition of metal oxides-

2Ag2O ⟶ 4Ag + O2

🡪 By the decomposition of hydrogen peroxide –

2H2O2 ⟶ 2H2O + O2

🡪 By electrolytic decomposition of water –

2H2O ⟶ 2H2 + O2

🡪 It is a colorless and odorless gas. 

🡪 It is soluble in water. 

🡪 Its freezing point is 55 K and melting point is 90 K.

🡪 It reacts with almost all metals and non-metals. 

🡪 Reaction with Ca metal and sulfur dioxide gas -

2Ca + O2 ⟶ 2CaO

2SO2 + O2V2O₅ ⟶ 2SO3

It is a vital gas for survival of human beings and many other organisms as it is used in respiration. 

Oxygen cylinders are used by mountaineers and in hospitals.

It is used in welding and combustion. 


Simple Oxides 


What are Oxides? 

An oxide is a chemical compound that contains at least one oxygen atom and one other element. Dianion of oxygen is also called oxide which is represented by O-2. All oxides compounds contain at least one dianion of oxygen. Oxides are generally binary compounds composed of oxygen and another element. 

Examples of Oxides – Al2O3 – Aluminium oxide, CO2 – Carbon dioxide, SO2 – Sulfur dioxide, CaO – Calcium oxide, MgO – Magnesium oxide, Na2O – Sodium oxide etc. 


Classification of Oxides 

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Oxides can be divided into following types on the basis of valency of another element in oxides –

  • Simple oxides 

  • Mixed oxides 

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Simple Oxides – Simple oxides are made up of one metal or semimetal and oxygen. These oxides carry only that number of oxygen atoms which is allowed by the normal valency of the element or metal. 

Examples of simple oxides – H2O, MgO, CaO, SiO2 etc. 

Mixed Oxides – Mixed oxides are produced when simple oxides combine. These two simple oxides can be of the same metal (element) or different. 

Examples of Mixed oxides – Red lead (Pb3O) is mixed oxide of lead dioxide (PbO2) and lead monoxide (PbO). Another example is ferro-ferric oxide (Fe3O4) which is a mixed oxide of two simple oxides – ferric oxide (Fe2O3) and ferrous oxide (FeO).


Oxides can be Divided into Following Types on the Basis of Metallic Character of another Element in Oxides –

Metallic oxides 

  • Basic oxide 

  • Amphoteric oxide 

Non-metallic oxide 

  • Acidic oxide 

  • Neutral oxide 

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Metallic Oxides – Metallic oxides are made of metal and oxygen. These are generally found in nature as minerals. These are formed by oxidation of metals.

Examples of metallic oxide – CaO, MgO, Fe3O4, BaO, ZnO etc. 

These can be classified into following two types –

  • Basic oxide 

  • Amphoteric oxide 


Basic Oxide – If an oxide reacts with water and forms a base is called basic oxide. Thus, basic oxide on reacting with water gives base. It means if we prepare a solution of basic oxide and water and dip a red litmus paper in it then it turns blue. 

Examples of basic oxides – MgO, CaO, BaO etc. 


Amphoteric Oxide – An amphoteric oxide is that metallic oxide which shows dual behavior. It behaves as an acidic oxide and basic oxide both. It also reacts with both bases as well as acids. 

Examples of Amphoteric oxides – Zinc oxide (ZnO)

When zinc oxide reacts with conc. Sodium hydroxide it acts as acidic oxide while when it reacts with HCl it acts as basic oxide. Reactions are given below –

ZnO + 2H2O + 2NaOH 🡪 Na3Zn[OH]4 + H2

                            Acidic zinc oxide

ZnO + 2HCl 🡪 ZnCl2 + H2O

                                  Basic zinc oxide 

Another example of amphoteric oxide is Al2O3 – aluminium oxide. When it reacts with sulfuric acid it acts as a base while when it reacts with sodium hydroxide it acts as acid. Reactions are given below –

Al2O3 + 3H2SO4 🡪 Al2(SO4)3 + 3H2O

                                             Basic

Al2O3 + 2NaOH 🡪 2NaAlO2 + H2O

                                              Acidic                      Sodium aluminate

Other examples of amphoteric oxides are BeO, SnO etc. 


Non- Metallic oxide – Non - Metallic oxides are formed by non - metal and oxygen. These are generally found in nature as gases such as carbon dioxide. These are formed by oxidation of non - metals. 

Examples of metallic oxide – CO2, SO2, P2O5, CO etc. 

These can be classified into following two types –

  • Acidic oxide 

  • Neutral oxide 


Acidic Oxides - If an oxide reacts with water and forms an acid is called acidic oxide. Thus, acidic oxide on reacting with water gives base. It means if we prepare a solution of acidic oxide and water and dip a blue litmus paper in it then it turns red. Mostly acidic oxides are oxides of non – metals but some oxides of metals with high oxidation states also possess acidic character. Thus, few metallic oxides such as CrO3, Mn2O7 etc. are also acidic oxides. 

Examples of acidic oxides – SO2, CO2, SO3 etc. 

When sulfur trioxide reacts with water, it forms sulfuric acid. Reaction is given below –

SO3 + H2O 🡪 H2SO4


Neutral Oxide – Neutral oxides are those oxides which neither show acidic properties nor basic properties. They do not form any salt when react with acid or base. 

Examples of neutral oxides – N2O, NO, CO etc. 


Ozone 

Ozone 

Preparation 

Properties 

Uses 

🡪 It is an allotropic form of oxygen. 

🡪 Its chemical formula is O3

🡪 It is found in the earth’s atmosphere (Stratosphere) in the form of a layer which is known as the ozone layer. 

🡪 Naturally, ozone is produced in the stratosphere by the photochemical decomposition of oxygen through UV rays. It takes place in two steps. In the 1st step oxygen molecule gets split into two oxygen atoms through UV rays. In the 2nd step, nascent oxygen reacts with other oxygen molecules and form ozone. 


Overall reaction –

3O2 Sunlight ⟶ 2O3

🡪 To convert oxygen into ozone, oxygen is passed through a silent electrical discharge. It is an endothermic process. 

Reaction –

3O2 ⟶ 2O3 (ΔH = +142 kJ/mol)


🡪 Pure ozone is of different colors in different states. In gaseous state, it is of pale blue color, in liquid state, it is of dark blue color and in solid state it is of violet – black color. 

🡪 Ozone is harmful in high concentrations. 

🡪 It is a powerful oxidizing agent. 

🡪 It combines with nitrogen oxides rapidly. 

Reaction – 

NO + O3 ⟶ NO2 + O2

🡪 It is used in ozone therapy which is used to treat many diseases caused by various bacteria, fungi and viruses. 

🡪 It is used in sterilizing water, disinfectants etc.

🡪 It is used as an oxidizing agent in production of KMnO4


Sulphur 

Following are the two important allotropes of Sulphur –

  • ɑ – Sulphur (Rhombic sulfur)

  • β – Sulphur (Monoclinic sulfur)

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Allotropic forms of Sulphur 

Appearance 

Properties 

Structure 

ɑ – Sulphur 

It is yellow in color. It is also called rhombic sulfur.  


🡪 It is the most commonly found form of Sulphur in nature. It is insoluble in water.

🡪 It is highly soluble in carbon disulfide and slightly soluble in benzene, alcohol and ether.  

🡪 Its melting point is 385.8 K. 

It contains S8 molecules which are puckered in a crown shape. It has rhombic crystal structure and octahedral shape. 

β – Sulphur 

It is also yellow in colored solid. It is also called monoclinic sulfur. 

🡪 Its melting point is 393 K. 

🡪 It is also soluble in carbon disulfide. 

It contains S8 molecules which are puckered in a crown shape. It is a monoclinic crystalline solid. It has needle shape. 

  • Monoclinic sulfur is stable at above 369 K temperature while rhombic sulfur is stable at below 369 K temperature. 


Sulphur Dioxide 

Sulphur Dioxide 

Preparation 

Properties 

Uses 

Chemical Formula – SO2 

🡪It is formed when sulfur is burnt in presence of oxygen. A small quantity of sulfur trioxide also forms with SO2

Reaction – 

S + O2 ⟶ SO2

🡪 Laboratory method – It is prepared in labs by treating sulphite with dilute sulphuric acid. 

🡪 Industrial method – At industrial level, it is prepared by roasting of sulphide ores. 

Reaction – 

4FeS2 + 11O2 ⟶ 2Fe2O3 + 8SO2

🡪 It is a colorless gas and has a pungent smell. 

🡪 It is highly soluble in water.

SO2 + H2O ⟶ H2SO3

🡪 Reaction with sodium hydroxide –

2NaOH + SO2 ⟶ Na2SO3 + H2

After forming Na2SO3, it reacts with excess of SO2 and forms sodium hydrogen sulphite. 

Reaction –

Na2SO3 + H2O + SO2 ⟶  2NaHSO3

🡪 In moist form, it acts as a reducing agent. 

 🡪 Sulphur dioxide reacts with chlorine in the presence of charcoal which acts as a catalyst and gives sulphuryl chloride.

Reaction - 

SO2(g) + Cl2(g) ⟶ SO2Cl2(l)  

🡪 Sulfur dioxide reacts with oxygen in presence of catalyst vanadium pentoxide and forms sulfur trioxide. 

Reaction – 

2SO2 + O2 ⟶ 2 SO3

🡪 It is used in refining petroleum and sugar. 

🡪 It is used in bleaching of cloths such as silk etc.

🡪 It is used as an antichlor, disinfectant and preservative. 

🡪 Liquid sulfur dioxide is used as a solvent to dissolve organic and inorganic chemicals. 


Oxoacids of Sulphur 

Sulphur forms many oxoacids. An oxoacid is an acid that contains oxygen. Oxoacids of sulfur contain oxygen and sulfur. For examples, H2SO3, H2S2O3, H2S2O7 etc. Different oxoacids of sulfur are prepared by different methods. The most commonly used and the most important oxoacid of sulphur is Sulfuric acid. 

Sulphuric acid is an acidic chemical compound with the formula H2SO4. It is widely used in the industries that is why it is called the king of chemicals. Its worldwide production clearly indicates its industrial strength. In the year 2004, its world production was about 180 million tonnes. It is also known as oil of vitriol. It is an odorless, colorless and viscous liquid which is soluble in water. Preparation of sulphuric acid generally involves highly exothermic processes.

Study of the oil of vitriol began in ancient times. It is believed that Muhammad ibn Zakariya al-Razi was the 1st alchemist of Iran who produced sulfuric acid. Then in 17th century German-Dutch Chemist Johann Glauber prepared sulfuric acid by sulfur with potassium nitrate. In 1736, Doctor Joshua Ward used this method for large scale production of sulfuric acid, although it was an expensive method of production. After this many other methods of production of sulfuric acid were also discovered but they were not economically feasible. Then in 1831, British vinegar merchant Peregrine Phillips patented the contact process, in which sulfuric acid is produced by using sulfur dioxide and oleum in presence of vanadium pentoxide as catalyst. This method is more economically feasible than available all other methods and produces concentrated sulfuric acid. It is the current method of producing sulfuric acid in large scale and high concentration required for industrial processes. 


Contact Process for Manufacturing of Sulfuric Acid 

Steps involved in contact process of manufacturing of sulfuric acid are listed below–

  • Preparation of sulfur dioxide 

  • Oxidation of sulfur dioxide to prepared sulfur trioxide 

  • Addition reaction of sulfur trioxide and sulfuric acid to give oleum 

  • Dilution of oleum to produce concentrated sulfuric acid 


Preparation of Sulfur Dioxide – In the 1st step sulfur is oxidized or burned to produce sulfur dioxide. Reaction is given below –

S(s) + O2(g) → SO2(g)


Oxidation of Sulfur Dioxide to Prepared Sulfur Trioxide – Sulfur dioxide is oxidized to sulfur trioxide in presence of vanadium pentoxide as catalyst. It is an exothermic reaction which is reversible in nature. the reaction is given below –

2SO2(g) + O2(g) V2O₅ ↔ 2SO3(g)


Addition Reaction of Sulfur Trioxide and Sulfuric Acid to Give Oleum – Sulfur trioxide is absorbed into ~98% sulfuric acid to form oleum which is also known as fuming sulfuric acid. It is an addition reaction. The reaction is given below –

SO3(g) + H2SO4(l) → H2S2O7(l)

                                 Oleum 


Dilution of Oleum to Produce Concentrated Sulfuric Acid – Oleum is diluted with water to form concentrated sulfuric acid. The reaction is given below –

H2S2O7(l) + H2O(l) → 2H2SO4(l)

It should be noticed here that we used 1 mole of sulfuric acid as reactant and produces 2moles of sulfuric acid. 

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