The P-Block Elements (Compounds Of Group 16 Elements)

Dioxygen

Dioxygen ($O_2$) is a vital element for life on Earth, playing a critical role in respiration and combustion.

Preparation

1. Laboratory Preparation:

Decomposition of Potassium Chlorate: Heating potassium chlorate ($KClO_3$) in the presence of a catalyst like manganese dioxide ($MnO_2$).

$2KClO_3(s) \xrightarrow[heat]{\Delta} 2KCl(s) + 3O_2(g)$

Heating Metal Peroxides: Heating metal peroxides like barium peroxide ($BaO_2$) in the presence of a catalyst.

$2BaO_2(s) \xrightarrow{catalyst} 2BaO(s) + O_2(g)$

Decomposition of Hydrogen Peroxide: Gentle heating of hydrogen peroxide ($H_2O_2$) in the presence of a negative catalyst like manganese dioxide ($MnO_2$) or potassium iodide ($KI$).

$2H_2O_2(aq) \xrightarrow{MnO_2 \ or \ KI} 2H_2O(l) + O_2(g)$

Action of Nitric Acid on Silver:

$2Ag(s) + 2HNO_3(dilute) \rightarrow 2AgNO_3(aq) + H_2O(l) + \frac{1}{2}O_2(g)$

2. Commercial Production:

Fractional Distillation of Liquid Air: This is the primary industrial method. Liquid air is fractionally distilled, separating nitrogen (boiling point -196°C) from oxygen (boiling point -183°C).
Electrolysis of Water: Electrolysis of acidified or alkaline water produces hydrogen at the cathode and oxygen at the anode.

$2H_2O(l) \xrightarrow{electrolysis} 2H_2(g) + O_2(g)$

Properties

Physical Properties:

Colorless, odorless, and tasteless gas.
Essential for respiration and combustion.
Slightly soluble in water.
Liquefies at 90 K (-183°C) and solidifies at 54 K (-219°C).
Liquid oxygen is pale blue.
Solid oxygen exists in several forms, the most stable being blue $\alpha$-oxygen.

Chemical Properties:

1. Combustion: Oxygen supports combustion, causing many substances to burn vigorously.

With Non-metals:

$C(s) + O_2(g) \rightarrow CO_2(g)$

$S(s) + O_2(g) \rightarrow SO_2(g)$

$4P(s) + 5O_2(g) \rightarrow P_4O_{10}(s)$

$2H_2(g) + O_2(g) \xrightarrow{catalyst} 2H_2O(g)$

With Metals: Reacts with most metals upon heating to form metal oxides.

$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

$4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3(s)$

With Compounds: Oxidizes compounds like $H_2S$, $NH_3$, $CH_4$, $C_2H_5OH$, etc.

$2H_2S(g) + 3O_2(g) \rightarrow 2SO_2(g) + 2H_2O(g)$

$4NH_3(g) + 5O_2(g) \xrightarrow{Pt} 4NO(g) + 6H_2O(g)$

$CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$

2. Oxidizing Agent: Oxygen is a strong oxidizing agent.

Its oxidizing property is due to its high electronegativity and tendency to form $O^{2-}$, $O_2^{2-}$, $O_2^-$, or $O_2$ (in ozone).
Oxidizes non-metals, metals, and even some compounds.

3. Formation of Oxides: Forms oxides with almost all elements.

4. Allotropes: Oxygen exists in two allotropic forms: dioxygen ($O_2$) and ozone ($O_3$).

Simple Oxides

Oxides: Oxides are binary compounds of oxygen with another element.

Classification Based on Chemical Properties:

Acidic Oxides: Formed by most non-metals. They react with water to form acids or react with bases to form salt and water.

Examples: $CO_2$ (forms $H_2CO_3$), $SO_2$ (forms $H_2SO_3$), $SO_3$ (forms $H_2SO_4$), $P_4O_{10}$ (forms $H_3PO_4$), $Cl_2O_7$ (forms $HClO_4$).

Basic Oxides: Formed by most metals. They react with water to form bases or react with acids to form salt and water.

Examples: $Na_2O$ (forms $NaOH$), $CaO$ (forms $Ca(OH)_2$), $MgO$ (forms $Mg(OH)_2$).

Amphoteric Oxides: Exhibit both acidic and basic properties. They react with both acids and bases.

Examples: $Al_2O_3$, $ZnO$, $PbO$, $SnO_2$, $BeO$, $As_2O_3$, $Sb_2O_3$.

$Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$

$Al_2O_3 + 2NaOH + 3H_2O \rightarrow 2Na[Al(OH)_4]$

Neutral Oxides: Do not react with either acids or bases.

Examples: $CO$, $N_2O$, $NO$, $H_2O$.

Trend in Group 16:

Oxygen oxides: $O_2$ is neutral, $O_3$ is a strong oxidizing agent, $H_2O_2$ is acidic and an oxidizer/reducer.
Sulfur oxides: $SO_2$ and $SO_3$ are acidic.
Selenium and Tellurium oxides ($SeO_2$, $TeO_2$) are amphoteric.
Polonium oxides ($PoO_2$, $PoO_3$) are basic.

This trend reflects the decrease in non-metallic character and increase in metallic character down the group.

Ozone

Ozone ($O_3$) is an allotrope of oxygen, consisting of three oxygen atoms bonded together.

Preparation

1. Silent Electric Discharge: Passing dry oxygen gas through a special electric discharge tube (ozonizer) converts oxygen into ozone.

$$3O_2(g) \xrightarrow{Electric \ discharge} 2O_3(g) \quad (\Delta H^\circ = +142.7 \text{ kJ/mol})$$

Properties:

This reaction is reversible and endothermic.
Ozone is a pale blue gas with a characteristic smell.
Liquid ozone is deep blue and solid ozone is violet-black.
Ozone is diamagnetic, while oxygen is paramagnetic.
It is thermodynamically unstable relative to oxygen and decomposes readily, especially in the presence of catalysts like $NO$, $Ag$, or $Pt$.

$2O_3(g) \rightarrow 3O_2(g) \quad (\Delta H^\circ = -142.7 \text{ kJ/mol})$

Strong Oxidizing Agent: Ozone is a much stronger oxidizing agent than dioxygen. It oxidizes many elements and compounds.

Oxidizes $NO$ to $NO_2$: $O_3 + NO \rightarrow O_2 + NO_2$
Oxidizes $H_2O_2$ to $O_2$: $O_3 + H_2O_2 \rightarrow 2O_2 + H_2O$
Bleaches colored substances by oxidation.

Bleaching Action: Due to its oxidizing nature.
Identification: Its presence can be detected by its characteristic smell or by its ability to turn starch-iodide paper blue-black (due to liberation of iodine from $KI$).

$2KI(aq) + O_3(g) + H_2O(l) \rightarrow 2KOH(aq) + I_2(aq) + O_2(g)$

Uses:

As a bleaching agent for cotton, oils, etc.
As a disinfectant for sterilizing water and purifying air.
In the synthesis of potassium periodate ($KIO_4$).
In the production of ink in modern printing process.

Ozone Layer: Atmospheric ozone ($O_3$) forms a layer in the stratosphere which absorbs harmful ultraviolet (UV) radiation from the sun, protecting life on Earth.

Sulphur — Allotropic Forms

Sulfur exhibits allotropy, existing in several different structural forms.

Common Allotropes:

Rhombic Sulfur ($\alpha$-Sulfur):

Structure: Crystals are orthorhombic. Consists of crown-shaped $S_8$ molecules.
Properties: Stable below 369 K (transition temperature). Yellow, crystalline solid.

Monoclinic Sulfur ($\beta$-Sulfur):

Structure: Crystals are monoclinic. Also consists of $S_8$ molecules, but arranged differently.
Properties: Stable above 369 K. Pale yellow, needle-shaped crystals.

Plastic Sulfur ($\gamma$-Sulfur):

Structure: Formed by rapid cooling of molten sulfur. Consists of long chains of sulfur atoms.
Properties: Amorphous, rubbery solid. Unstable and slowly converts to rhombic sulfur.

Other forms: At higher temperatures, sulfur exists as $S_2$ molecules, similar to oxygen.

Transition Temperature: The temperature at which one allotrope converts into another (369 K for $\alpha$-S to $\beta$-S) is called the transition temperature.

Sulphur Dioxide

Sulfur dioxide ($SO_2$) is a colorless gas with a pungent, suffocating odor.

Preparation

1. Laboratory Preparation:

From Sulfites/Bisulfites and Acids: Heating a sulfite or bisulfite salt with a strong non-oxidizing acid.

$Na_2SO_3(s) + H_2SO_4(dilute) \xrightarrow{heat} Na_2SO_4(aq) + H_2O(l) + SO_2(g)$

$2NaHSO_3(s) + H_2SO_4(dilute) \rightarrow Na_2SO_4(aq) + 2H_2O(l) + 2SO_2(g)$

From Sulfites and HCl:

$Na_2SO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + SO_2(g)$

2. Commercial Production:

Roasting of Sulfides: The primary industrial method involves heating metal sulfides (like iron pyrites, $FeS_2$) in excess air.

$4FeS_2(s) + 11O_2(g) \xrightarrow{heat} 2Fe_2O_3(s) + 8SO_2(g)$

Burning of Sulfur: Burning sulfur in air or oxygen.

$S(s) + O_2(g) \rightarrow SO_2(g)$

Properties

Physical Properties:

Colorless gas.
Pungent, suffocating smell.
Denser than air.
Liquefies at 263 K (-10°C) and solidifies at 197.7 K (-75.4°C).
Readily soluble in water, forming sulfurous acid ($H_2SO_3$).

$SO_2(g) + H_2O(l) \rightleftharpoons H_2SO_3(aq)$

Chemical Properties:

1. Reducing Agent: $SO_2$ acts as a reducing agent, especially in acidic solutions. It gets oxidized to sulfate ($SO_4^{2-}$).

Reduces acidified potassium permanganate ($KMnO_4$).

$5SO_2(g) + 2KMnO_4(aq) + 2H_2O(l) \rightarrow K_2SO_4(aq) + 2MnSO_4(aq) + 2H_2SO_4(aq)$

Reduces dichromate ions ($Cr_2O_7^{2-}$).

$3SO_2(g) + K_2Cr_2O_7(aq) + H_2SO_4(aq) \rightarrow K_2SO_4(aq) + Cr_2(SO_4)_3(aq) + H_2O(l)$

Bleaches colored substances by reducing them.

2. Oxidizing Agent: $SO_2$ can also act as an oxidizing agent, particularly in reactions where it is reduced to sulfide ($S^{2-}$).

Reduces $H_2S$ to Sulfur.

$SO_2(g) + 2H_2S(g) \rightarrow 3S(s) + 2H_2O(l)$

3. Acidic Nature: It is an acidic oxide.

Reacts with water to form sulfurous acid ($H_2SO_3$), a weak acid.
Reacts with basic oxides and hydroxides to form sulfites.

$SO_2(g) + 2NaOH(aq) \rightarrow Na_2SO_3(aq) + H_2O(l)$

4. Combination with Oxygen: Catalytically combines with oxygen to form sulfur trioxide ($SO_3$), the key step in the Contact process for sulfuric acid production.

$2SO_2(g) + O_2(g) \xrightarrow{V_2O_5 \ catalyst} 2SO_3(g)$

Uses:

Manufacture of sulfuric acid.
As a bleach for wool, silk, and paper pulp (reducing bleach).
As a preservative for fruits and vegetables (by preventing oxidation).
As a solvent in liquefaction processes.
As a fumigating agent.

Pollution: $SO_2$ is a major air pollutant contributing to acid rain.

Oxoacids Of Sulphur

Sulfur forms a variety of oxoacids, characterized by the presence of sulfur atoms bonded to oxygen atoms, often with S-O bonds and sometimes S-S or S-H bonds.

Common Oxoacids:

Sulfur Dioxide Oxoacids:

Sulfurous Acid ($H_2SO_3$): Formed by the reaction of $SO_2$ with water. It is a weak acid and a reducing agent.

$SO_2(g) + H_2O(l) \rightleftharpoons H_2SO_3(aq)$

It forms sulfites ($SO_3^{2-}$) and bisulfites ($HSO_3^-$).

Sulfur Trioxide Oxoacids:

Sulfuric Acid ($H_2SO_4$): The most important oxoacid, discussed separately.
Sulfuric Trioxide ($SO_3$): Acts as the anhydride of sulfuric acid.

$SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)$

Other Oxoacids: Sulfur forms a series of oxoacids with sulfur in various oxidation states. These include:

Hyposulfurous Acid ($H_2S_2O_4$): Contains sulfur in the $+3$ oxidation state. It is a strong reducing agent.
Dithionous Acid ($H_2S_2O_4$): Similar to hyposulfurous acid.
Thiosulfurous Acid ($H_2SO_2$): Less common.
Disulfurous Acid ($H_2S_2O_5$): Exists in equilibrium with sulfurous acid and sulfur dioxide.
Pyrosulfurous Acid ($H_2S_2O_5$): Not a true acid but exists in equilibrium.
Dithionic Acid ($H_2S_2O_6$): Contains sulfur in $+5$ oxidation state.
Peroxodisulfuric Acid ($H_2S_2O_8$): Contains a peroxide linkage ($O-O$). Sulfur is in $+6$ oxidation state. Strong acid and oxidizing agent.

$H_2S_2O_8 + 2H_2O \rightarrow 2H_2SO_4 + H_2O_2$

Monothionic Acid ($H_2SO_4$): Sulfuric acid itself.
Polythionic Acids ($H_2S_nO_6$): A series of acids with a chain of sulfur atoms.

Structure: Oxoacids of sulfur generally contain $S=O$ bonds and $S-OH$ groups. The presence of P-H bonds in phosphorus oxoacids is absent in sulfur oxoacids.

Sulphuric Acid

Sulfuric acid ($H_2SO_4$) is one of the most important industrial chemicals, often referred to as the "king of chemicals".

Manufacture (Contact Process)

Overall Reaction: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$

Steps:

Production of Sulfur Dioxide ($SO_2$): By burning sulfur or roasting sulfide ores in excess air.

$S(s) + O_2(g) \rightarrow SO_2(g)$

$4FeS_2(s) + 11O_2(g) \rightarrow 2Fe_2O_3(s) + 8SO_2(g)$

Catalytic Oxidation of $SO_2$ to $SO_3$: This is the crucial step and is carried out in the presence of a catalyst in the Contact Process.

$2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) \quad (\Delta H^\circ = -196.6 \text{ kJ/mol})$

Catalyst: Vanadium(V) oxide ($V_2O_5$) supported on silica ($SiO_2$).
Conditions: Temperature around 720 K (450°C) and pressure of 1-2 atm. These conditions are chosen to maximize the yield of $SO_3$ according to Le Chatelier's principle (exothermic reaction favored by low temperature, but high temperature needed for reasonable rate).

Absorption of $SO_3$ in Sulfuric Acid: Sulfur trioxide ($SO_3$) is absorbed in concentrated sulfuric acid (98%) to form oleum ($H_2S_2O_7$). This is done to prevent the formation of a fine mist of sulfuric acid, which is difficult to condense.

$SO_3(g) + H_2SO_4(l) \rightarrow H_2S_2O_7(l)$

Dilution of Oleum: Oleum is then diluted with water to get sulfuric acid of desired concentration.

$H_2S_2O_7(l) + H_2O(l) \rightarrow 2H_2SO_4(l)$

Properties

Physical Properties:

Pure sulfuric acid is a colorless, odorless, viscous liquid.
It is hygroscopic (absorbs moisture from the air) and miscible with water in all proportions with evolution of heat.
Boiling point: 611 K (338°C). It decomposes on boiling.
High density ($\approx 1.84$ g/mL).
Strong dehydrating agent.

Chemical Properties:

1. Acidic Nature: Sulfuric acid is a strong dibasic acid, ionizing in two steps:

Step 1: $H_2SO_4 + H_2O \rightarrow H_3O^+ + HSO_4^-$ (strong acid)
Step 2: $HSO_4^- + H_2O \rightleftharpoons H_3O^+ + SO_4^{2-}$ (weak acid)
Forms acidic ($HSO_4^-$) and normal ($SO_4^{2-}$) salts.

2. Dehydrating Agent:

It removes water from other compounds. It chars carbohydrates like sugar and paper by abstracting hydrogen and oxygen in the ratio 2:1.

$C_{12}H_{22}O_{11}(s) \xrightarrow{conc. H_2SO_4} 12C + 11H_2O$

Used to dry gases like $H_2$, $O_2$, $CO_2$, $SO_2$ (but not ammonia or hydrogen iodide, which react with it).

3. Oxidizing Agent: Concentrated sulfuric acid is a strong oxidizing agent, especially when hot.

With Metals: Oxidizes metals like copper, silver, lead, etc., to their sulfates, getting reduced to $SO_2$.

$Cu(s) + 2H_2SO_4(conc.) \xrightarrow{heat} CuSO_4(aq) + SO_2(g) + 2H_2O(l)$

With Non-metals: Oxidizes non-metals like carbon and sulfur.

$C(s) + 2H_2SO_4(conc.) \xrightarrow{heat} CO_2(g) + 2SO_2(g) + 2H_2O(l)$

$S(s) + 2H_2SO_4(conc.) \xrightarrow{heat} 3SO_2(g) + 2H_2O(l)$

With Halides: Concentrated $H_2SO_4$ oxidizes $HI$ and $HBr$ to $I_2$ and $Br_2$, respectively.

$2HI + H_2SO_4(conc.) \rightarrow I_2(s) + SO_2(g) + 2H_2O(l)$

4. Action with Salts:

Reacts with salts of more volatile acids to produce those acids (displacement).

$NaCl(s) + H_2SO_4(conc.) \rightarrow NaHSO_4(s) + HCl(g)$

$NaNO_3(s) + H_2SO_4(conc.) \rightarrow NaHSO_4(s) + HNO_3(g)$

Uses:

Manufacture of fertilizers (e.g., superphosphates).
In the production of sulfuric acid itself, detergents, dyes, pigments, explosives, and other chemicals.
As a dehydrating agent in laboratories and industries.
As an oxidizing and dehydrating agent in organic synthesis.
In lead storage batteries.