Why are Group 16 elements called chalcogens?

The term 'chalcogen' originates from Greek words: 'khalkos' meaning 'ore' and 'genes' meaning 'producing' or 'forming'. This name was given because many of these elements, particularly oxygen and sulfur, are frequently found in the Earth's crust as components of various ores. For instance, many metal oxides (like iron ore) and sulfides (like galena, $PbS$) are common minerals. This historical observation of their prevalence in ore formation led to their collective designation as chalcogens.

Explain the anomalous behavior of oxygen.

Oxygen, as the first member of Group 16, exhibits properties significantly different from the rest of the group. This is primarily due to its exceptionally small size, very high electronegativity (second only to fluorine), and the complete absence of vacant d-orbitals in its valence shell. These factors prevent oxygen from expanding its octet, limiting its maximum covalency to four (e.g., in $H_3O^+$) and typically two. Its high electronegativity also enables strong hydrogen bonding, leading to water being a liquid while other hydrides ($H_2S, H_2Se$) are gases.

How do the acidic and reducing characters of Group 16 hydrides change down the group?

Both the acidic character and the reducing character of Group 16 hydrides ($H_2E$) increase as you move down the group. For acidic character, as the size of the central atom (E) increases, the E-H bond length increases, and its strength decreases. This makes it easier to break the E-H bond and release a proton ($H^+$), thus increasing acidity ($H_2O < H_2S < H_2Se < H_2Te$). Similarly, the decreasing bond strength and thermal stability down the group mean these hydrides can more readily donate hydrogen or electrons, making them stronger reducing agents ($H_2O < H_2S < H_2Se < H_2Te$). Water is practically non-reducing, while $H_2Te$ is a strong reducing agent.

Why is $SF_6$ exceptionally stable, while $SCl_6$ does not exist?

$SF_6$ is remarkably stable due to two main reasons. Firstly, sulfur can achieve a $+6$ oxidation state by expanding its octet using its vacant 3d-orbitals, forming six strong S-F bonds. Secondly, and more importantly, the six small fluorine atoms sterically protect the central sulfur atom very effectively. This 'steric hindrance' prevents attacking reagents (like water) from approaching the sulfur atom, making $SF_6$ kinetically inert. In contrast, chlorine atoms are much larger than fluorine atoms. Six chlorine atoms would be too bulky to fit around the sulfur atom without significant steric repulsion, making $SCl_6$ unstable and thus it does not form.

What is the inert pair effect and how does it apply to Group 16 elements?

The inert pair effect refers to the reluctance of the $ns^2$ valence electrons to participate in bonding in heavier p-block elements. As we move down Group 16, the stability of the $+6$ oxidation state decreases, while the stability of the $+4$ oxidation state increases. This is because the $ns^2$ electrons become more tightly held by the nucleus due to poor shielding by intervening d- and f-electrons, making them less available for bonding. Consequently, only the $np^4$ electrons participate, leading to a $+4$ oxidation state. For example, for tellurium, $Te(IV)$ compounds are more stable than $Te(VI)$ compounds, and for polonium, $Po(II)$ and $Po(IV)$ are more common than $Po(VI)$.

Why is the electron gain enthalpy of oxygen less negative than that of sulfur?

Electron gain enthalpy is generally expected to become less negative down a group due to increasing atomic size. However, oxygen is an exception. Despite its position, its electron gain enthalpy is less negative (less exothermic) than that of sulfur. This anomaly arises because oxygen has a very small atomic size. When an incoming electron approaches the compact $2p$ subshell of oxygen, it experiences significant inter-electronic repulsion from the already present electrons. This repulsion makes the addition of an electron less favorable, resulting in a less negative (or less exothermic) electron gain enthalpy compared to sulfur, which has a larger atomic size and less electron density in its $3p$ subshell, allowing for easier electron accommodation.

Group 16 Elements

Q: Why is $SF_6$ exceptionally stable, while $SCl_6$ does not exist?

$SF_6$ is remarkably stable due to two main reasons. Firstly, sulfur can achieve a $+6$ oxidation state by expanding its octet using its vacant 3d-orbitals, forming six strong S-F bonds. Secondly, and more importantly, the six small fluorine atoms sterically protect the central sulfur atom very effectively. This 'steric hindrance' prevents attacking reagents (like water) from approaching the sulfur atom, making $SF_6$ kinetically inert. In contrast, chlorine atoms are much larger than fluorine atoms. Six chlorine atoms would be too bulky to fit around the sulfur atom without significant steric repulsion, making $SCl_6$ unstable and thus it does not form.

Q: What is the inert pair effect and how does it apply to Group 16 elements?

The inert pair effect refers to the reluctance of the $ns^2$ valence electrons to participate in bonding in heavier p-block elements. As we move down Group 16, the stability of the $+6$ oxidation state decreases, while the stability of the $+4$ oxidation state increases. This is because the $ns^2$ electrons become more tightly held by the nucleus due to poor shielding by intervening d- and f-electrons, making them less available for bonding. Consequently, only the $np^4$ electrons participate, leading to a $+4$ oxidation state. For example, for tellurium, $Te(IV)$ compounds are more stable than $Te(VI)$ compounds, and for polonium, $Po(II)$ and $Po(IV)$ are more common than $Po(VI)$.

Q: Why is the electron gain enthalpy of oxygen less negative than that of sulfur?

Electron gain enthalpy is generally expected to become less negative down a group due to increasing atomic size. However, oxygen is an exception. Despite its position, its electron gain enthalpy is less negative (less exothermic) than that of sulfur. This anomaly arises because oxygen has a very small atomic size. When an incoming electron approaches the compact $2p$ subshell of oxygen, it experiences significant inter-electronic repulsion from the already present electrons. This repulsion makes the addition of an electron less favorable, resulting in a less negative (or less exothermic) electron gain enthalpy compared to sulfur, which has a larger atomic size and less electron density in its $3p$ subshell, allowing for easier electron accommodation.

Chemistry

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Version 1Updated 22 Mar 2026

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Group 16 elements, often referred to as the chalcogens, constitute the second group of p-block elements in the modern periodic table. This group comprises oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po). They are characterized by having six valence electrons in their outermost shell, specifically an $ns^2 np^4$ electronic configuration. This configuration drives their chem…

Quick Summary

Group 16 elements, known as chalcogens, include Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), and Polonium (Po). They all possess an $ns^2 np^4$ valence electron configuration, aiming to gain two electrons to achieve a stable octet, thus commonly exhibiting a $-2$ oxidation state.

However, elements from sulfur onwards can also show $+2, +4$ , and $+6$ oxidation states due to the availability of vacant d-orbitals. Key trends down the group include increasing atomic size, decreasing ionization enthalpy and electronegativity, and a transition from non-metallic to metallic character.

Oxygen displays anomalous behavior due to its small size, high electronegativity, and absence of d-orbitals, leading to unique properties like hydrogen bonding in water. Allotropes are common, such as $O_2$ and $O_3$ for oxygen, and rhombic and monoclinic sulfur.

Their hydrides ( $H_2E$ ) show increasing acidic and reducing character but decreasing thermal stability and bond angles down the group. Halides vary in stability and oxidation states, with $SF_6$ being notably stable due to steric protection.

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Key Concepts

Electronic Configuration and Oxidation States

The defining characteristic of Group 16 elements is their $ns^2 np^4$ valence electron configuration. This…

Trends in Hydrides (

H_2E

)

Group 16 elements form hydrides of the general formula $H_2E$ . The properties of these hydrides exhibit clear…

Anomalous Behavior of Oxygen

Oxygen stands apart from the other chalcogens due to its unique characteristics. Its exceptionally small…

Elements — O, S, Se, Te, Po (Chalcogens)

Valence e- config —

ns^2 np^4

Common O.S. —

-2

(O),

-2, +2, +4, +6

(S, Se, Te, Po)

Trends Down Group

- Atomic/Ionic size: Increases - Ionization Enthalpy: Decreases - Electronegativity: Decreases - Metallic Character: Increases - Acidic character of $H_2E$ : Increases ( $H_2O < H_2S < H_2Se < H_2Te$ ) - Thermal stability of $H_2E$ : Decreases ( $H_2O > H_2S > H_2Se > H_2Te$ ) - Reducing character of $H_2E$ : Increases ( $H_2O < H_2S < H_2Se < H_2Te$ ) - Bond angle in $H_2E$ : Decreases ( $H_2O (104.5^circ) > H_2S (92.1^circ)$ )

Anomalous O — Small size, high EN, no d-orbitals. Leads to H-bonding in

H_2O

, max covalency 4, less negative electron gain enthalpy than S.

Allotropes — O (

O_2, O_3

), S (Rhombic

S_8

, Monoclinic

S_8

, Plastic).

Key Compounds —

OF_2

(O is

+2

SF_6

(stable, octahedral),

SF_4

(see-saw).

To remember Group 16 elements: Old Students Select Tellurium Polonium. (Oxygen, Sulfur, Selenium, Tellurium, Polonium)