The Chalcogens: Chemistry Of Group 16 Elements

The Group 16 elements, also known as the chalcogens, are a family of elements positioned in Group 16 (formerly Group VIA) of the periodic table. The elements of this group are oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po) which is radioactive. The chalcogens derive their name from the Greek word “chalkos,” meaning copper, as they are commonly found in copper ores.

Table of Contents

Chemistry of Group 16 Elements

Generally, this group of elements exhibit several common characteristics because of their similar electronic configurations and chemical properties. They are all nonmetals, except for polonium, which is a metalloid at the p-block of the periodic table. The chalcogens are known for their diverse roles in various chemical reactions and biological processes.

Occurrence and Extraction of Group 16 Elements

Natural Abundance and Distribution of the Chalcogens

Oxygen (O)

Oxygen is the most abundant element in the Earth’s crust, constituting approximately 46% of its mass. Asides its molecular form, it is found in combined state such as oxides, silicates, and carbonates, which are prevalent in rocks, minerals, and the atmosphere.

Sulfur (S)

Sulfur is widely distributed in nature and occurs in various minerals, such as sulfides (e.g., pyrite), sulfates (e.g., gypsum), and elemental sulfur deposits. It is present in sedimentary rocks, volcanic deposits, and fossil fuels. Sulfur is also released into the atmosphere during volcanic eruptions and industrial processes.

Selenium (Se)

Selenium occurs in trace amounts in the Earth’s crust. It is often associated with sulfur-containing minerals, such as pyrite and chalcopyrite, as well as in certain organic deposits and sedimentary rocks. Selenium can be found in soil, water bodies, and plants.

Tellurium (Te)

Tellurium is a relatively rare element in the Earth’s crust which also occurs in trace amounts. It is primarily found in telluride minerals associated with gold, silver, and copper deposits. Tellurium is also present in some coal deposits and as a byproduct of copper refining.

Polonium (Po)

Polonium is a highly radioactive element and occurs in extremely trace amounts in the Earth’s crust. It is primarily obtained as a decay product of uranium and thorium.

Find out about The Occurence of Alkali Earth Metals (Group 2 Elements) and more

Extraction Methods of Group 16 Elements

Oxygen (O)

Oxygen is not extracted directly from ores or minerals, however, it is obtained from the atmosphere through air separation techniques, such as cryogenic distillation or pressure swing adsorption.

The cryogenic distillation which is the most common method for oxygen extraction involves cooling air to extremely low temperatures (-183°C) to liquefy it. The liquefied air is then separated into its components using fractional distillation. Oxygen, having a lower boiling point than nitrogen, vaporizes first and is collected.

Sulfur (S)

Sulfur extraction involves mining sulfur-rich deposits, including underground mining of sulfur ores, such as pyrite and galena. The mined ores undergo crushing and grinding processes to release elemental sulfur. The extracted sulfur is then refined using various techniques, including the Frasch process (involving superheated water) or the Claus process (involving hydrogen sulfide gas).

Frasch Process: This process is suitable for extracting sulfur from underground deposits. High-pressure superheated water (steam) is injected into the deposit, which melts the sulfur and forms a liquid sulfur-oil mixture. The molten sulfur is then pumped to the surface using compressed air.

Claus Process: The Claus process is used for extracting sulfur from natural gas and petroleum refining. In this process, the sulfur-containing gases, such as hydrogen sulfide (H₂S), are burned to produce sulfur dioxide (SO₂). The SO₂ gas is then reacted with more H₂S in the presence of catalysts to form elemental sulfur.

2H₂S + SO₂ -> 3S + 2H₂O

Selenium (Se)

Selenium is often extracted as a byproduct during the extraction and refining of other metals, such as copper, lead, and nickel. The process involves smelting the ores containing selenium, followed by separation techniques to isolate selenium compounds from the impurities.

The copper anode slimes undergo further refining processes to separate selenium from other elements. One common method involves roasting the slimes, which converts selenium into selenium dioxide (SeO₂). The SeO₂ can then be leached and purified to obtain selenium metal or selenium compounds.

Tellurium (Te)

Tellurium is obtained as a byproduct of copper refining processes. The ores containing tellurium are smelted, and the resulting copper anode slimes are further processed to separate tellurium from other elements.

Polonium (Po)

Polonium is primarily produced through neutron irradiation of bismuth-209 in specialized nuclear reactors. The irradiated bismuth is then chemically processed to isolate polonium.

Industrial Applications

Oxygen (O)

Oxygen has variety of industrial applications, including combustion support in industries, steelmaking, metal refining, wastewater treatment, and medical applications such as oxygen therapy and anesthesia.

Sulfur (S)

Sulfur is used in the manufacturing of sulfuric acid, which is turn used to produce fertilizers, detergents, pharmaceuticals, and dyes. Sulfur is also utilized in rubber vulcanization, petroleum refining, and as a component in batteries.

Selenium (Se)

Selenium has applications in electronics, glass manufacturing, and photovoltaics. Its compounds are also utilized in pigments, glass coloration, and pharmaceuticals.

Tellurium (Te)

Tellurium has applications in various industries, including solar energy, electronics, thermoelectric materials, and alloys. It is a crucial component in high-efficiency solar cells, as well as in rewritable optical discs and semiconductors.

Polonium (Po)

Polonium though has limited application due to its radioactive nature however, it is used in anti-static devices, as a heat source in thermoelectric power generators, and in nuclear research applications.

Do you know the applications of Group 8 Elements?

Physical Properties and Appearances of Group 16 Elements (Chalcogens)

General Physical Characteristics of Group 16 Elements

The chalcogens are nonmetals, except for polonium, which is a metalloid or post-transition metal.

These elements exhibit a wide range of physical and chemical properties due to variations in atomic structure and bonding characteristics.

Physical States and Melting/Boiling Points of Group 16 Element

Group 16 Element	Physical State	Melting Point (°C)	Boiling Point (°C)	Allotropes	Density (g/cm³)	Atomic Radius (pm)	Color and Appearance
Oxygen (O)	Gas	-218.79°C	-182.96°C	O₂ (oxygen), O₃ (ozone)	1.429 g/cm³	152 pm	Colorless gas (liquid and solid forms appear pale blue)
Sulfur (S)	Solid	115.21°C	444.61°C	S₈ rings (α-sulfur, plastic sulfur, etc.)	2.07 g/cm³	180 pm	Yellow crystals
Selenium (Se)	Solid	217°C	685°C	Amorphous selenium, red selenium, gray selenium	4.81 g/cm³	190 pm	Gray metallic appearance
Tellurium (Te)	Solid	449.5°C	988°C	Amorphous tellurium, gray tellurium, black tellurium	6.24 g/cm³	206 pm	Silvery-white or grayish metalloid with a metallic luster
Polonium (Po)	Solid	254°C	962°C	Cubic crystal structure	9.196 g/cm³	197 pm	Silvery-gray metal with high reflectivity

The physical states of Group 16 elements vary due to differences in intermolecular forces and bonding types.

Oxygen is a diatomic gas (O₂) with a strong double bond.

Sulfur, selenium, and tellurium are solids at room temperature due to stronger intermolecular forces resulting from larger atomic sizes and the presence of multiple valence electrons.

Polonium is also a solid, but its melting and boiling points are relatively low due to the weak metallic bonding in its crystal structure.

Allotropes of Group 16 Elements

Group 16 elements exhibit allotropy due to variations in the arrangement of atoms or molecules in different crystalline or amorphous structures.

The different allotropes arise from differences in bonding, packing arrangements, and molecular structures, which can lead to variations in physical properties such as color and stability.

Oxygen (O)

Molecular Oxygen (O₂): This is the most common and stable form of oxygen. It consists of a diatomic molecule with two oxygen atoms bonded together by a double bond. Molecular oxygen is a colorless, odorless gas that is essential for supporting combustion and sustaining life.

Oxygen is highly reactive and supports combustion, making it essential for respiration and many chemical processes. Under normal conditions, O₂ is relatively unreactive, but it readily participates in various reactions, such as combustion, oxidation, and reduction reactions. For example, oxygen reacts with most elements to form oxides, like the reaction of oxygen with iron to form iron oxide (rust):

4Fe + 3O₂ -> 2Fe₂O₃

Ozone (O₃): Ozone is a less stable form of oxygen. It consists of three oxygen atoms bonded together. Ozone is a pale blue gas with a distinct odor. It plays a vital role in the Earth’s atmosphere, where it absorbs ultraviolet (UV) radiation, protecting life on Earth. It is also a powerful oxidizing agent. It exhibits higher reactivity than molecular oxygen. Ozone readily decomposes to molecular oxygen and nascent oxygen, which is highly reactive. It participates in reactions like ozonolysis, where it breaks carbon-carbon double bonds. An example is the reaction of ozone with an alkene:

CH₂=CH₂ + O₃ -> CH₂O + CH₂O + O₂

Sulfur (S)

Rhombic Sulfur (α-S): This is the most common and stable form of sulfur at room temperature. It consists of S₈ rings, where eight sulfur atoms are arranged in a puckered ring. Rhombic sulfur is yellow and crystalline in appearance. It is insoluble in water and has a low melting point.

It has low electrical conductivity and is an insulator. It undergoes combustion to form sulfur dioxide:

S + O₂ -> SO₂

Rhombic sulfur also reacts with metals to form metal sulfides, such as the reaction with iron to form iron sulfide:

Fe + S -> FeS

Monoclinic Sulfur (β-S): Monoclinic sulfur is a metastable form of sulfur. It consists of S₈ rings, similar to rhombic sulfur, but with a different arrangement. Monoclinic sulfur is also yellow, but it is denser and has a higher melting point than rhombic sulfur. It exhibits similar chemical reactivity to rhombic sulfur.

Selenium (Se)

Amorphous Selenium: Amorphous selenium is a non-crystalline form of selenium. It is formed when molten selenium is rapidly cooled or when selenium vapor is condensed rapidly. Amorphous selenium is dark brown or black and lacks the crystalline structure of other allotropes. It exhibits photoconductivity, meaning its electrical conductivity increases upon exposure to light. It can be converted to crystalline forms by heating. It reacts with halogens to form selenium halides, like the reaction with chlorine to form selenium dichloride:

Se + Cl₂ -> SeCl₂

Hexagonal Selenium: Hexagonal selenium is the most stable crystalline form of selenium. It consists of a layered structure where each layer contains helical chains of selenium atoms. It is grayish-black and exhibits semiconductor properties. It reacts with strong acids to form selenic acid:

Se + H₂SO₄ -> H₂SeO₄

Tellurium (Te)

α-Tellurium: α-Tellurium is the most common and stable form of tellurium. It has a trigonal crystal structure and is silvery-white in color. It is a semiconductor material and exhibits metallic conductivity when heated. α-Tellurium can be brittle and is a good thermal conductor. It reacts with acids to form tellurium dioxide, such as the reaction with nitric acid:

Te + 4HNO₃ -> TeO₂ + 4NO₂ + 2H₂O

β-Tellurium: β-Tellurium is a metastable form of tellurium. It has a distorted, orthorhombic crystal structure. It is darker in color than α-tellurium and exhibits metallic properties.

Polonium (Po)

α-Polonium: α-Polonium is the most stable and commonly occurring form of polonium. It has a simple cubic crystal structure and is silvery in appearance. α-Polonium undergoes radioactive decay, emitting alpha particles.

β-Polonium: β-Polonium is a metastable form of polonium. It has a body-centered cubic crystal structure and is slightly denser than α-polonium.

See the Allotropes in the Carbon Group

Atomic Properties of Chalcogens

Appearance

Descending the group, the appearance of the Chalcogens changes. Oxygen is a colorless gas, sulfur is a yellow solid, selenium is a gray solid, tellurium is a silvery-white solid, and polonium is a silvery solid. The trend is towards darker colors as we go down the group.

The change in appearance is primarily due to the increasing atomic size and the consequent variation in electronic structure and energy levels, which influence the absorption and reflection of light.

Electronic Configuration

The electronic configurations of the Chalcogens follow a pattern as we descend the group. The general electronic configuration is ns²np⁴, where n represents the principal quantum number.

The principal quantum number (n) increases as we move down the group, indicating the addition of new energy levels. This results in a more extended electronic configuration with additional occupied orbitals.

Atomic Radius

The atomic radius generally increases as we descend the Chalcogen group.

The increase in atomic radius is primarily due to the addition of new energy levels (principal quantum number) as we move down the group. Each successive element has an extra electron shell, leading to an increase in the size of the electron cloud surrounding the nucleus.

Ionization Energy

Ionization energy generally decreases as we move down the group of Chalcogens.

The decreasing trend is primarily attributed to the increasing atomic size. As the atomic radius increases, the valence electrons are further from the nucleus, resulting in weaker attraction between the nucleus and the outermost electrons. Consequently, less energy is required to remove an electron, leading to lower ionization energies.

Electronegativity

Electronegativity generally decreases as we descend the Chalcogen group.

The decreasing trend is mainly due to the increasing atomic size. Larger atoms have more electron shielding from the nucleus, which reduces the effective nuclear charge experienced by the valence electrons. As a result, the attraction for electrons decreases, leading to lower electronegativity values.

Electron Affinity

Electron affinity, the energy change when an atom gains an electron, does not exhibit a consistent trend down the Chalcogen group.

The electron affinity values can vary due to the complex interplay between factors such as atomic size, electron shielding, and electron configuration. These factors can lead to variations in the stability of the electron configurations of the resulting anions.

Chalcogen	Appearance	Electronic Configuration	Atomic Radius	Ionization Energy	Electronegativity	Electron Affinity
Oxygen (O)	Colorless Gas	1s² 2s² 2p⁴	73 pm	1314 kJ/mol	3.44	140 kJ/mol
Sulfur (S)	Yellow Solid	1s² 2s² 2p⁶ 3s² 3p⁴	104 pm	999 kJ/mol	2.58	200 kJ/mol
Selenium (Se)	Gray Solid	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴	116 pm	941 kJ/mol	2.55	195 kJ/mol
Tellurium (Te)	Silvery-White Solid	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁴	123 pm	869 kJ/mol	2.1	190 kJ/mol
Polonium (Po)	Silvery Solid	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁴	135 pm	812 kJ/mol	2.0	183 kJ/mol

Oxidation States of Chalcogens

Chalcogen	Oxidation States	Examples
Oxygen (O)	-2	H₂O (Water), CO₂ (Carbon Dioxide)
Sulfur (S)	-2, -1, +4, +6	H₂S (Hydrogen Sulfide), S^2- (Sulfide ion), SO₂ (Sulfur Dioxide), H₂SO₄ (Sulfuric Acid)
Selenium (Se)	-2, -1, +4, +6	H₂Se (Hydrogen Selenide), Se^2- (Selenide ion), SeO₂ (Selenium Dioxide), H₂SeO₄ (Selenic Acid)
Tellurium (Te)	-2, -1, +4, +6	H₂Te (Hydrogen Telluride), Te^2- (Telluride ion), TeO₂ (Tellurium Dioxide), H₂TeO₄ (Telluric Acid)
Polonium (Po)	-2, -1, +2, +4, +6	Po^2- (Polonide ion), PoO₂ (Polonium Dioxide), H₂PoO₃ (Polonic Acid)

Chemical Properties of Group 16 Elements

Oxidation-Reduction Reactions

Oxygen (O)

Oxygen is highly reactive and readily participates in oxidation-reduction reactions. For example, oxygen can react with hydrogen (H₂) to form water (H₂O) through a redox reaction:

2H₂ + O₂ → 2H₂O

In this reaction, oxygen is reduced from an oxidation state of 0 to -2, while hydrogen is oxidized from an oxidation state of 0 to +1.

Sulfur (S)

Similarly, Sulfur can react with oxygen (O₂) to form sulfur dioxide (SO₂):

S + O₂ → SO₂

In this reaction, sulfur is oxidized from an oxidation state of 0 to +4, while oxygen is reduced from an oxidation state of 0 to -2.

Selenium (Se)

Also, Selenium can undergo oxidation-reduction reactions as it reacts with chlorine (F₂) to form selenium hexafluoride (SeF₆):

Se + 3F₂ → SeF₆

In this reaction, selenium is oxidized from an oxidation state of -2 to +6, while chlorine is reduced from an oxidation state of 0 to -1.

Tellurium (Te)

In the same vein, Tellurium can react with hydrogen peroxide (H₂O₂) to form tellurium dioxide (TeO₂):

Te + H₂O₂ → TeO₂ + H₂O

In this reaction, tellurium is oxidized from an oxidation state of 0 to +4, while hydrogen peroxide is reduced from an oxidation state of 0 to -2.

Polonium (Po)

Likewise, the radioactive Polonium can react with oxygen (O₂) to form polonium dioxide (PoO₂):

Po + O₂ → PoO₂

Where polonium is oxidized from an oxidation state of 0 to +4, while oxygen is reduced from an oxidation state of 0 to -2.

The Chalcogens: Chemistry of Group 16 Elements

Acid-Base Reactions

Oxygen (O)

Oxygen can act as a weak acid in certain reactions, but it is not typically involved in direct acid-base reactions like sulfur, selenium, tellurium, and polonium.

Sulfur (S)

Sulfur can act as both an acid and a base in different reactions. For example, sulfuric acid (H₂SO₄) is a strong acid and can donate hydrogen ions (H⁺) in aqueous solutions:

H₂SO₄ → 2H⁺ + SO₄²⁻

In this reaction, sulfuric acid donates two protons (H⁺) to the solution, forming sulfate ions (SO₄²⁻).

Selenium (Se)

Selenium can also exhibit acidic and basic properties in certain reactions. For example, Selenium dioxide (SeO₂) can act as an acid when it reacts with a base, such as sodium hydroxide (NaOH):

SeO₂ + 2NaOH → Na₂SeO₃ + H₂O

In this reaction, selenium dioxide acts as an acid by donating a proton (H⁺) to sodium hydroxide, forming sodium selenite (Na₂SeO₃) and water (H₂O).

Tellurium (Te)

Similarly, Tellurium can behave as an acid or a base depending on the reaction conditions.

For example, Telluric acid (H₂TeO₄) can act as an acid in aqueous solutions:

H₂TeO₄ → 2H⁺ + TeO₄²⁻

In this reaction, telluric acid donates two protons (H⁺) to the solution, forming tellurate ions (TeO₄²⁻).

Polonium (Po)

Polonium, being a radioactive element, is less commonly involved in direct acid-base reactions compared to sulfur, selenium, and tellurium.

Polonium can form polonides (anions) when it reacts with certain metals:

Po + 2Na → Na₂Po

In this reaction, polonium reacts with sodium (Na) to form sodium polonide (Na₂Po).

Formation of Chalcogenides

Oxygen (O)

Oxygen can form chalcogenides with certain elements, typically referred to as oxides rather than chalcogenides. Oxides are compounds where oxygen is bonded to another element.

For instance, Oxygen can react with metals to form metal oxides, such as iron oxide (Fe₂O₃):

4Fe + 3O₂ → 2Fe₂O₃

In this reaction, iron (Fe) reacts with oxygen (O₂) to form iron(III) oxide (Fe₂O₃).

Sulfur (S)

Sulfur readily forms chalcogenides with metals and nonmetals, commonly known as metal sulfides and nonmetal sulfides, respectively.

For example, Sulfur reacts with metals to form metal sulfides, such as iron sulfide (FeS):

Fe + S → FeS

In this reaction, iron (Fe) reacts with sulfur (S) to form iron sulfide (FeS).

Selenium (Se)

Similarly, Selenium can also form chalcogenides with metals and nonmetals. Infact, when Selenium reacts with metals, it forms metal selenides, such as copper selenide (Cu₂Se):

Cu + Se → Cu₂Se

Tellurium (Te)

Same happens with Tellurium which reacts with metals to form metal tellurides, such as lead telluride (PbTe):

Pb + Te → PbTe

Polonium (Po)

However, Polonium is less commonly involved in the formation of chalcogenides due to its rarity and radioactive nature. However, it can still form compounds known as polonides.

For example, Polonium reacts with certain metals to form metal polonides, such as bismuth polonide (BiPo):

Bi + Po → BiPo

Reducing Agents

Oxygen (O)

Oxygen is not typically considered a strong reducing agent. It is more commonly involved in oxidation reactions as an oxidizing agent.

Sulfur (S)

Sulfur can act as a reducing agent in certain reactions, especially when it forms sulfide compounds.

Hydrogen sulfide (H₂S) can reduce certain metal ions, such as silver ions (Ag⁺):

H₂S + 2Ag⁺ → 2H⁺ + 2Ag + S

In this reaction, hydrogen sulfide (H₂S) donates electrons to silver ions (Ag⁺), reducing them to silver (Ag) while sulfur (S) is formed.

Selenium (Se)

Selenium can also exhibit reducing properties, especially when it forms selenide compounds.

Selenium can reduce certain metal ions, such as mercury(II) ions (Hg²⁺):

Hg²⁺ + Se → HgSe

In this reaction, selenium (Se) donates electrons to mercury(II) ions (Hg²⁺), reducing them to form mercury selenide (HgSe).

Tellurium (Te)

Similarly, tellurium can act as a reducing agent, particularly when it forms telluride compounds.

For example, tellurium can reduce certain metal ions, such as gold(III) ions (Au³⁺):

Au³⁺ + 2Te → AuTe₂

In this reaction, tellurium (Te) donates electrons to gold(III) ions (Au³⁺), reducing them to form gold telluride (AuTe₂).

Polonium (Po)

However, Polonium is a radioactive element and less commonly involved in reducing reactions compared to sulfur, selenium, and tellurium.

It has been proven that polonium can reduce certain metal ions, such as lead(IV) ions (Pb⁴⁺):

Pb⁴⁺ + Po → Pb²⁺ + Po²⁺

In this reaction, polonium (Po) donates electrons to lead(IV) ions (Pb⁴⁺), reducing them to lead(II) ions (Pb²⁺) while polonium is oxidized to polonium(II) ions (Po²⁺).

Formation of Oxoacids

Chalcogens can form oxoacids, which contain oxygen and hydrogen.

Oxygen (O)

Oxygen is not typically involved in the direct formation of oxoacids as it is already present in many common oxoacids such as sulfuric acid (H₂SO₄) and selenic acid (H₂SeO₄).

Sulfur (S)

Sulfur can form various oxoacids, with sulfuric acid being the most well-known and commonly used oxoacid.

For instance, Sulfur can form sulfuric acid (H₂SO₄) through a series of oxidation steps:

S + O₂ → SO₂

2SO₂ + O₂ → 2SO₃

SO₃ + H₂O → H₂SO₄

In this reaction sequence, sulfur is successively oxidized to sulfur dioxide (SO₂) and then to sulfur trioxide (SO₃), which reacts with water (H₂O) to produce sulfuric acid (H₂SO₄).

Selenium (Se)

Selenium can also form oxoacids, with selenic acid being one of the notable examples.

Selenium can form selenic acid (H₂SeO₄), an oxoacids through oxidation reactions:

Se + 2H₂O + 3O₂ → H₂SeO₄

In this reaction, selenium (Se) reacts with water (H₂O) and oxygen (O₂) to form selenic acid (H₂SeO₄).

Tellurium (Te)

Similarly, tellurium can form telluric acid (H₆TeO₆) through oxidation reactions:

Te + 4H₂O + 6O₂ → H₆TeO₆

In this reaction, tellurium (Te) reacts with water (H₂O) and oxygen (O₂) to form telluric acid (H₆TeO₆).

5. Polonium (Po):

Though Polonium has limited practical applications in the formation of oxoacids. However, it can potentially form polonic acid (H₂PoO₄).

Po + 2H₂O + 4O₂ → H₂PoO₄

In this reaction, polonium (Po) reacts with water (H₂O) and oxygen (O₂) to form polonic acid (H₂PoO₄).

Formation of Polyatomic Anions

Oxygen (O)

Oxygen can form the oxide ion (O²⁻) by gaining two electrons.:

O + 2e⁻ → O²⁻

Sulfur (S)

Same can be said for Sulfur which forms the sulfide ion (S²⁻) by gaining two electrons.

S + 2e⁻ → S²⁻

Selenium (Se)

Selenium can form various polyatomic anions, such as selenide ion (Se²⁻) and selenate ion (SeO₄²⁻), by gaining electrons.

Se + 2e⁻ → Se²⁻

Tellurium (Te)

Tellurium can form polyatomic anions, including telluride ion (Te²⁻) and tellurate ion (TeO₄²⁻), by gaining electrons:

Te + 2e⁻ → Te²⁻

Polonium (Po)

Polonium can also potentially form polyatomic anions by gaining electrons, although its reactivity is relatively low compared to other Chalcogens. For example:

Polonium (Po) gains two electrons to form a hypothetical polonide ion (Po²⁻):

Po + 2e⁻ → Po²⁻

Some Exceptional Chemical Properties and Reactions

Oxygen (O)

Combustion: Oxygen is a key component for combustion reactions. For example, oxygen reacts with organic compounds in the presence of heat to produce carbon dioxide and water vapor.

C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O

Oxidizing Agent: Oxygen acts as a strong oxidizing agent in many reactions. It readily accepts electrons from other substances, leading to oxidation.

Formation of Peroxides: Oxygen can also form peroxides, such as hydrogen peroxide (H₂O₂), which is a strong oxidizing and bleaching agent.

Oxygen is highly reactive and readily participates in oxidation reactions. However, noble gases like helium and neon are unreactive towards oxygen, and no significant reactions occur.

Oxygen shows exceptional reactivity with certain transition metals, such as palladium (Pd) and platinum (Pt), forming metal oxides that exhibit catalytic properties. For example:

Pd + O₂ → PdO

Pt + O₂ → PtO₂

Sulfur (S)

Formation of Sulfur Oxides: Sulfur reacts with oxygen to form different sulfur oxides, such as sulfur dioxide (SO₂) and sulfur trioxide (SO₃). These compounds have various industrial applications.

Acid-Base Reactions: Sulfuric acid (H₂SO₄), derived from sulfur, is a highly corrosive and widely used strong acid.

Precipitation of Metal Sulfides: Sulfur reacts with metal ions to form insoluble metal sulfides, which often precipitate out of solution. This property is utilized in analytical chemistry for detecting metal ions.

While sulfur is generally known for its reactivity, certain forms of elemental sulfur exhibit low reactivity. For instance, sulfur in its stable rhombic form (S₈) is relatively unreactive at room temperature.

Sulfur can form complex reaction products with certain metals, such as mercury (Hg). One example is the formation of mercury sulfide (HgS):

Hg + S → HgS

Selenium (Se)

Redox Reactions: Selenium can undergo redox reactions, often exhibiting multiple oxidation states (-2, -1, +4, +6) similar to sulfur.

Formation of Selenides: Selenium can react with metals to form metal selenides, which have semiconducting properties.

Selenium’s reactivity is lower compared to sulfur and oxygen. It requires more rigorous conditions or reactants to undergo reactions.

Selenium can form unusual selenium-rich compounds with metals, such as the formation of selenides. However, some metals, like silver (Ag), do not readily react with selenium.

Selenium exhibits photochemical reactivity under appropriate conditions, involving energy transfer and excitation processes.

Tellurium (Te)

Tellurium has lower reactivity compared to sulfur and oxygen, similar to selenium.

Tellurium can form alloys with various metals, but its reactivity with some metals, such as gold (Au) and platinum (Pt), is limited.

Tellurium shows limited reactivity with nonmetals, such as halogens, compared to sulfur and selenium.

Semiconductor Applications: Tellurium is used in the manufacturing of semiconductors, especially in alloys with other elements like cadmium, mercury, and copper.

Polonium (Po)

Certainly! Here are some exceptions for the reactivity of each member of the Group 16 elements (Chalcogens), along with relevant chemical reactions:

Polonium is highly radioactive, and its reactivity is influenced by its radioactive decay.

Polonium exhibits limited reactivity due to its short half-life and scarcity in nature.

Polonium can undergo self-decay through alpha particle emission, transforming into a different element.

Chemistry of Group 16 Quiz

Compounds of Chalcogens

Oxides

Oxygen readily forms oxides with various elements. These can be prepared by the direct combination of oxygen with metals, nonmetals, or other compounds. For example:

2Mg + O₂ → 2MgO (Formation of magnesium oxide)

Metal Oxides

Metal oxides are compounds composed of a metal element bonded to oxygen. They can be prepared through various methods, including the direct combination of a metal with oxygen or by the thermal decomposition of metal compounds.

Metal oxides can exhibit a range of physical properties and appearances. For example:

Iron oxide (Fe₂O₃) occurs naturally as the reddish-brown mineral known as rust.

Titanium dioxide (TiO₂) is a white solid commonly used as a pigment in paints and cosmetics.

Metal oxides involve in

Acid-Base Reactions:

Metal oxides can react with acids to form salts and water.

For example: MgO + 2HCl → MgCl₂ + H₂O (Reaction between magnesium oxide and hydrochloric acid)

Reduction Reactions:

Metal oxides can be reduced by reacting with carbon monoxide (CO) or hydrogen (H₂) to produce the pure metal. For example:

Fe₂O₃ + 3CO → 2Fe + 3CO₂ (Reduction of iron oxide with carbon monoxide)

Non-metal Oxides

Nonmetal oxides are compounds composed of a nonmetal element bonded to oxygen. They are typically formed through the direct combination of a nonmetal with oxygen or the oxidation of nonmetallic compounds.

Nonmetal oxides can vary in physical properties and appearances.

For example: Carbon dioxide (CO₂) is a colorless gas at standard conditions.

Sulfur dioxide (SO₂) is a colorless gas with a pungent odor.

– Nonmetal oxides also involve:

Acid-Base Reactions

Nonmetal oxides can react with water to form acids. For example:

CO₂ + H₂O → H₂CO₃ (Reaction between carbon dioxide and water, forming carbonic acid)

CO₂ + CaO → CaCO₃

Redox Reactions: Nonmetal oxides can participate in redox reactions, either as oxidizing agents or reducing agents. For example:

SO₂ + 2H₂S → 3S + 2H₂O (Sulfur dioxide oxidizing hydrogen sulfide to elemental sulfur)

Amphoteric Oxides

Amphoteric oxides are compounds that exhibit both acidic and basic properties. They can act as either acids or bases, depending on the reaction conditions.

Amphoteric oxides include compounds like aluminum oxide (Al₂O₃) and zinc oxide (ZnO).

The chemical reactivity of amphoteric oxides depends on the reaction medium. They can react with both acids and bases, forming salts and water. For example:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (Reaction between aluminum oxide and hydrochloric acid)

Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (Reaction between aluminum oxide and sodium hydroxide)

Acidic Oxides

Acidic oxides are compounds that react with water to form acids. They typically contain nonmetal elements and have a high affinity for water.

Acidic oxides can be prepared by the reaction of nonmetallic elements with oxygen or through the oxidation of nonmetallic compounds.

Examples of acidic oxides include sulfur trioxide (SO₃) and nitrogen dioxide (NO₂).

When these oxides react with water, they form acids through the process of hydration. For example:

SO₃ + H₂O → H₂SO₄ (Sulfur trioxide reacts with water to form sulfuric acid)

Basic Oxides

Basic oxides, also known as basic anhydrides, are compounds that react with water to form bases.

Basic oxides are usually metal oxides, such as calcium oxide (CaO) and sodium oxide (Na₂O).

When these oxides react with water, they undergo hydrolysis to produce metal hydroxides, which are bases. For example:

CaO + H₂O → Ca(OH)₂ (Calcium oxide reacts with water to form calcium hydroxide)

Neutral Oxides

Neutral oxides neither exhibit acidic nor basic properties when they react with water.

Examples of neutral oxides are Nitrogen (I) oxide (N₂O), nitrogen (II) oxide (NO), carbon monoxide (CO). It does not undergo significant hydrolysis and does not produce acids or bases when in contact with water.

Peroxides

Peroxides, such as hydrogen peroxide (H₂O₂), can be prepared through various methods. One common method involves the reaction of a metal with hydrogen peroxide, while another method involves the reaction of an alkali metal superoxide with water. These reactions typically require appropriate conditions such as the presence of a catalyst or specific temperatures.

Appearance & Physical Properties

Hydrogen peroxide (H₂O₂) is a colorless liquid at room temperature with a slightly bitter taste. It has a pale blue color when it is highly concentrated.

Hydrogen peroxide has a relatively high boiling point of 150.2 °C (302.4 °F) and can decompose upon exposure to light or high temperatures.

It is miscible with water, alcohol, and ether, and has a higher density than water.

Hydrogen peroxide is a powerful oxidizing agent and can cause bleaching or whitening effects on various materials.

Chemical Reactivity of Peroxides

Hydrogen peroxide exhibits strong oxidizing properties and can readily donate oxygen atoms. It can react with various substances, including metals, nonmetals, and organic compounds.

It acts as a powerful bleaching agent due to its ability to oxidize pigments and break down organic compounds.

Hydrogen peroxide can react with certain metals, such as iron or manganese, in the presence of an acid catalyst, producing metal salts and water. For example:

2H₂O₂ + 2H⁺ + 2Fe²⁺ → 2Fe³⁺ + 4H₂O (Reaction between hydrogen peroxide, acid, and ferrous ions)

It can also act as a reducing agent in the presence of oxidizing agents, undergoing decomposition to release oxygen. For example:

2H₂O₂ → 2H₂O + O₂ (Decomposition of hydrogen peroxide to water and oxygen)

Superoxides

Superoxides are compounds that contain the superoxide ion (O₂⁻), which is a negatively charged oxygen species. They are typically formed by the reaction of alkali metals with excess oxygen or the reaction of alkaline earth metals with oxygen under specific conditions.

For example, potassium superoxide (KO₂) can be prepared by exposing potassium metal to a stream of pure oxygen:

4K + O₂ → 2KO₂

Superoxides are typically solid compounds with a characteristic yellow or orange color.

They have a high reactivity towards moisture and carbon dioxide in the air, which can lead to the formation of corresponding metal hydroxides and carbonates.

Superoxides are paramagnetic, meaning they are attracted to a magnetic field due to the presence of unpaired electrons.

Superoxides are powerful oxidizing agents and can readily donate oxygen atoms. They exhibit a higher reactivity compared to normal oxides.

– They can react with water, releasing oxygen and forming metal hydroxides. For example:

2KO₂ + 2H₂O → 2KOH + H₂O₂ + O₂

Superoxides also react with carbon dioxide, forming metal carbonates and releasing oxygen. For example: 2KO₂ + 2CO₂ → K₂CO₃ + O₂

– Superoxides can undergo thermal decomposition at elevated temperatures, releasing oxygen gas. This property makes them useful as a source of oxygen in certain applications.

– Due to their high reactivity, superoxides find applications in chemical synthesis, oxygen generation, and as components in self-contained breathing apparatus for emergency situations.

Oxoacids

Oxoacids are acids that contain oxygen and hydrogen, along with one or more other elements. They are typically prepared by the combination of oxygen with other elements, followed by the addition of hydrogen ions (protons) to the resulting compound.

The specific method of preparation varies depending on the oxoacid. Common methods include the reaction of an appropriate oxide or hydroxide with water, or the reaction of a corresponding anhydride with water.

CO₂ + H₂O à H₂CO₃

Appearance & Physical Properties

Oxoacids can exist in various physical forms, including liquids and solids, depending on the specific acid.

Sulfuric acid (H₂SO₄), for example, is a viscous, oily liquid that is colorless or slightly yellow in its pure form. It is highly corrosive and has a strong odor.

Phosphoric acid (H₃PO₄) is a white crystalline solid that is highly soluble in water.

Chemical Reactivity

Oxoacids are known for their acidic properties and their ability to donate protons (H⁺) in aqueous solutions.

They react with bases to form salts and water through neutralization reactions.

Oxoacids can undergo various other chemical reactions, such as oxidation or reduction reactions, depending on the specific acid and the conditions.

They can participate in acid-catalyzed reactions, serve as reactants in synthesis processes, or act as components in pH regulation and electrolyte solutions.

Examples of Oxoacids

Sulfuric acid (H₂SO₄): Prepared by the reaction of sulfur trioxide (SO₃) with water.

SO₃ + H₂O → H₂SO₄

Phosphoric acid (H₃PO₄): Prepared by the reaction of phosphorus pentoxide (P₂O₅) with water.

P₂O₅ + 3H₂O → 2H₃PO₄

Nitric acid (HNO₃): Prepared by the reaction of nitrogen dioxide (NO₂) with water.

3NO₂ + H₂O → 2HNO₃ + NO

Carbonic acid (H₂CO₃): Formed when carbon dioxide (CO₂) dissolves in water.

CO₂ + H₂O ⇌ H₂CO₃

Sulfides

Sulfur readily forms sulfides with metals and nonmetals. Sulfides can be prepared through direct combination reactions or by the reaction of hydrogen sulfide (H₂S) with metals.

For example, iron sulfide (FeS) is formed by the direct combination of iron (Fe) and sulfur (S):

Fe + S → FeS

Sulfides often exhibit distinct colors. Iron sulfide, for instance, appears black.

Sulfoxides and Sulfones

Sulfur can undergo oxidation to form sulfoxides and sulfones. These compounds contain a sulfur atom bonded to one or more oxygen atoms.

Dimethyl sulfoxide (DMSO) is a well-known example of a sulfoxide, with the formula (CH₃)₂SO. It has a liquid form and is a highly polar solvent.

Sulfuric Acid

Sulfuric acid (H₂SO₄) is a highly important compound used in various industrial processes, such as the production of fertilizers, dyes, detergents, and pharmaceuticals.

It can be prepared by the reaction of sulfur trioxide (SO₃) with water:

SO₃ + H₂O → H₂SO₄

Sulfuric acid is a dense, oily liquid that is colorless to slightly yellow. It is highly corrosive and has a strong odor. It is a strong acid and is widely used in laboratory and industrial settings.

In terms of chemical reactivity:

– Sulfides often exhibit low reactivity with acids and are commonly used as ores for metal extraction.

– Sulfoxides and sulfones have diverse applications, including as solvents, reagents in organic synthesis, and pharmaceutical intermediates.

– Sulfuric acid is a powerful dehydrating agent, capable of removing water from substances. It also exhibits strong acidic properties and can react with various compounds.

Selenides

Selenium can form selenides with metals and nonmetals. Selenides can be prepared through direct combination reactions. For example:

Cu + Se → CuSe (Formation of copper selenide)

Selenides often exhibit dark colors, such as copper selenide (CuSe) being brownish-black.

Selenoxides and Selenones

Selenium compounds with oxygen, such as selenoxides and selenones, can be formed through oxidation reactions.

Diselenides

Selenium can also form diselenides, which are compounds with a selenium-selenium bond. They can be prepared by reacting selenium with alkali metals, among other methods.

Tellurides

Tellurium forms tellurides with metals and nonmetals. Tellurides can be prepared through direct combination reactions. For example:

Ba + Te → BaTe (Formation of barium telluride)

Tellurides often have a metallic appearance and can be semiconductors.

Tellurites and Tellurates

Tellurium can form tellurites and tellurates, compounds that contain tellurium and oxygen. These compounds can be obtained through various oxidation processes.

Metalloid Tellurides

Tellurium can also form metalloid tellurides, which are compounds that exhibit both metallic and nonmetallic properties. They have applications in semiconductor devices.