DAT General Chemistry: Chalcogens and Allotropes of Oxygen

DAT High-Yield Concepts

Group 16 Valence Configuration: Chalcogens possess an $ns^2 np^4$ valence electron configuration (6 valence electrons) and typically achieve octet stability by gaining two electrons to form $-2$ anions or sharing electron pairs in covalent bonds.
Oxygen Allotropes: Dioxygen ( $\ce{O2}$ ) is paramagnetic with two unpaired electrons in $\pi^*$ antibonding orbitals. Ozone ( $\ce{O3}$ ) is a bent, diamagnetic allotrope with resonance-stabilized bonds and a bond angle of $116.8^\circ$ .
Sulfur Oxidation States: Unlike oxygen (which is constrained to $-2$ in most compounds due to high electronegativity), sulfur expands its octet using available d-orbitals, exhibiting oxidation states ranging from $-2$ in $\ce{H2S}$ to $+4$ in $\ce{SO2}$ and $+6$ in $\ce{H2SO4}$ .

1. The Chalcogen Family (Group 16)

The chalcogens (Group 16) consist of oxygen ( $\ce{O}$ ), sulfur ( $\ce{S}$ ), selenium ( $\ce{Se}$ ), tellurium ( $\ce{Te}$ ), polonium ( $\ce{Po}$ ), and livermorium ( $\ce{Lv}$ ). The term "chalcogen" originates from Greek words meaning "ore-forming," as many metal ores occur naturally as oxides or sulfides. Moving down Group 16, metallic character, atomic radius, and density increase, whereas electronegativity and ionization energy decrease. Unlike the adjacent halogen family (Group 17, covered in depth in its own lesson), which needs only one electron to complete an octet, chalcogens require two electrons, forming $-2$ oxidation states in binary ionic compounds with electropositive metals.

Assortment of chalcogen group elements. (Source: de:User:Tomihahndorf . · CC BY-SA 3.0 · Wikimedia Commons)

Mnemonic: Old Style Sewing Technique Polishes (

\ce{O}

\ce{S}

\ce{Se}

\ce{Te}

\ce{Po}

)

2. Allotropes of Oxygen: Dioxygen and Ozone

Oxygen exists primarily in two allotropic gas forms: molecular dioxygen ( $\ce{O2}$ ) and ozone ( $\ce{O3}$ ).

Allotropy: The property of an element to exist in two or more distinct physical forms with different structural arrangements and chemical properties.

Molecular dioxygen ( $\ce{O2}$ ) is a colorless, odorless gas essential for cellular respiration. Molecular orbital theory shows that $\ce{O2}$ possesses two unpaired electrons in degenerate $\pi^*_{2p}$ antibonding orbitals, explaining its paramagnetic character. In contrast, ozone ( $\ce{O3}$ ) is a reactive, pale blue gas with a pungent odor. Ozone is generated photochemically or by electrical discharge via the endothermic equilibrium:

\ce{3O2(g) <=> 2O3(g)}

The $\ce{O3}$ molecule is bent with $C_{2v}$ symmetry, a bond angle of $116.8^\circ$ , and resonance structures yielding a bond order of 1.5 with identical bond lengths of $\pu{127.8 pm}$ .

Property	Dioxygen ( $\ce{O2}$ )	Ozone ( $\ce{O3}$ )
Molecular Geometry	Linear	Bent ( $\approx 117^\circ$ )
Magnetism	Paramagnetic (2 unpaired $e^-$ )	Diamagnetic
Bond Order	2.0	1.5 (resonance)
Reactivity / Oxidizing Power	Moderate oxidant	Strong oxidant

3. Sulfur and Variable Oxidation States

Sulfur ( $\ce{S}$ ) is the second member of Group 16. In its elemental standard state, sulfur forms yellow, crown-shaped $\ce{S8}$ rings (cyclo-octasulfur). Due to lower electronegativity than oxygen and available valence 3d orbitals, sulfur exhibits diverse oxidation states ranging from $-2$ to $+6$ .

Frost diagrams (oxidation state vs. standard potential) for sulfur in (A) anoxic and (B) oxic solutions. The solid line signifies a solution pH = 0, while the dashed line represents a solution pH = 14. (Source: Иван Олефиренко · CC BY-SA 4.0 · Wikimedia Commons)

In aqueous solution, the relative thermodynamic stability of sulfur species across these oxidation states depends heavily on pH, as summarized by Frost diagrams. Under acidic conditions ( $\text{pH} = 0$ ), hydrogen sulfide ( $\ce{H2S}$ ) represents the most stable thermodynamic sink, whereas under basic conditions ( $\text{pH} = 14$ ), sulfate ( $\ce{SO4^{2-}}$ ) is the most thermodynamically favored species.

Worked Example: Assigning Sulfur Oxidation States

Problem: Determine the oxidation state of sulfur in hydrogen sulfide ( $\ce{H2S}$ ), elemental sulfur ( $\ce{S8}$ ), sulfur dioxide ( $\ce{SO2}$ ), and sulfuric acid ( $\ce{H2SO4}$ ).

Step-by-step Solution:

$\ce{H2S}$ : Hydrogen is assigned $+1$ . Summing charges: $2(+1) + x = 0 \implies x = -2$ .
$\ce{S8}$ : Free elemental forms have an oxidation state of $0$ .
$\ce{SO2}$ : Oxygen is assigned $-2$ . Summing charges: $x + 2(-2) = 0 \implies x = +4$ .
$\ce{H2SO4}$ : Hydrogen is $+1$ , oxygen is $-2$ . Summing charges: $2(+1) + x + 4(-2) = 0 \implies +2 + x - 8 = 0 \implies x = +6$ .

DAT Exam Trap: On the DAT, do not confuse oxidation state with formal charge! While formal charge assumes equal sharing of bonding electrons, oxidation state assigns shared electrons entirely to the more electronegative atom. Also, remember that elemental allotropes like

\ce{S8}

\ce{O3}

always have an average oxidation state of

0

Critical Thinking Questions

Why does oxygen max out at an oxidation state of $+2$ (in $\ce{OF2}$ ), whereas sulfur can reach a $+6$ oxidation state in $\ce{SF6}$ or $\ce{H2SO4}$ ?
How does Molecular Orbital Theory account for the paramagnetism of $\ce{O2}$ despite having an even number of valence electrons?
Using Frost diagram principles, explain why sulfite ( $\ce{SO3^{2-}}$ , oxidation state $+4$ ) can act as both an oxidizing agent and a reducing agent in solution.

References

Sources returned by Google Search grounding during lesson verification.

wikipedia.org — vertexaisearch.cloud.google.com
ebsco.com — vertexaisearch.cloud.google.com
sciencenotes.org — vertexaisearch.cloud.google.com
uleth.ca — vertexaisearch.cloud.google.com
study.com — vertexaisearch.cloud.google.com
echemi.com — vertexaisearch.cloud.google.com
libretexts.org — vertexaisearch.cloud.google.com
libretexts.org — vertexaisearch.cloud.google.com
stackexchange.com — vertexaisearch.cloud.google.com
wikipedia.org — vertexaisearch.cloud.google.com
Chalcogen - Wikipedia — en.wikipedia.org
7.4: Frost Diagrams - Chemistry LibreTexts — chem.libretexts.org

Chalcogens and Allotropes of Oxygen

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