Important Compounds of Sulfur

Introduction to the Major Compounds of Sulfur

Sulfur (S), a p-block element belonging to Group 16, forms a variety of compounds exhibiting diverse oxidation states ranging from -2 to +6. Its compounds are pivotal in industrial processes, environmental chemistry, and biological systems. For high school chemistry, the most frequently tested compounds include its oxides (like Sulfur Dioxide), oxyacids (like Sulfuric Acid), and hydrides (like Hydrogen Sulfide).

Sulfur Dioxide ($\text{SO}_2$)

Chemical Formula

$\text{SO}_2$

Common Name

Sulfur dioxide

Laboratory Preparation

Sulfur dioxide is typically prepared in the laboratory by the action of dilute sulfuric acid on a sulfite or bisulfite salt.

Balanced Equation: $\text{Na}_2\text{SO}_3 (\text{s}) + \text{H}_2\text{SO}_4 (\text{dil}) \rightarrow \text{Na}_2\text{SO}_4 (\text{aq}) + \text{H}_2\text{O} (\text{l}) + \text{SO}_2 (\text{g})$

Alternatively, using sodium bisulfite: $\text{NaHSO}_3 (\text{s}) + \text{HCl} (\text{dil}) \rightarrow \text{NaCl} (\text{aq}) + \text{H}_2\text{O} (\text{l}) + \text{SO}_2 (\text{g})$

Industrial Preparation (Brief)

Industrially, $\text{SO}_2$ is produced by roasting sulfide ores, such as iron pyrites ($\text{FeS}_2$), or by burning sulfur.

Balanced Equation: $\text{4FeS}_2 (\text{s}) + \text{11O}_2 (\text{g}) \rightarrow \text{2Fe}_2\text{O}_3 (\text{s}) + \text{8SO}_2 (\text{g})$

Sulfuric Acid ($\text{H}_2\text{SO}_4$)

Chemical Formula

$\text{H}_2\text{SO}_4$

Common Names

Oil of Vitriol, King of Chemicals

Preparation Process: Contact Process

Sulfuric acid is manufactured by the Contact Process, which involves several steps:

1. Production of Sulfur Dioxide ($\text{SO}_2$): By burning sulfur in air or roasting sulfide ores. $\text{S} (\text{s}) + \text{O}_2 (\text{g}) \rightarrow \text{SO}_2 (\text{g})$ or $\text{4FeS}_2 (\text{s}) + \text{11O}_2 (\text{g}) \rightarrow \text{2Fe}_2\text{O}_3 (\text{s}) + \text{8SO}_2 (\text{g})$

2. Catalytic Oxidation of $\text{SO}_2$ to $\text{SO}_3$: Sulfur dioxide is oxidized to sulfur trioxide ($\text{SO}_3$) using a vanadium pentoxide ($\text{V}_2\text{O}_5$) catalyst at optimal conditions (450-500°C and 1.5-1.7 atm pressure). This step is reversible and exothermic. $\text{2SO}_2 (\text{g}) + \text{O}_2 (\text{g}) \xrightarrow{\text{V}_2\text{O}_5, 450-500^\circ\text{C}, 1.5-1.7 \text{ atm}} \text{2SO}_3 (\text{g})$

3. Absorption of $\text{SO}_3$ in Sulfuric Acid: To prevent the formation of acid mist, $\text{SO}_3$ is absorbed in concentrated $\text{H}_2\text{SO}_4$ to form oleum (fuming sulfuric acid, $\text{H}_2\text{S}_2\text{O}_7$). $\text{SO}_3 (\text{g}) + \text{H}_2\text{SO}_4 (\text{conc}) \rightarrow \text{H}_2\text{S}_2\text{O}_7 (\text{l})$

4. Dilution of Oleum: Oleum is then diluted with a calculated amount of water to obtain sulfuric acid of desired concentration. $\text{H}_2\text{S}_2\text{O}_7 (\text{l}) + \text{H}_2\text{O} (\text{l}) \rightarrow \text{2H}_2\text{SO}_4 (\text{aq})$

Properties and Exam-Relevant Reactions

Physical Properties

• Colorless, dense, oily, viscous liquid.
• Highly corrosive.
• Strong affinity for water (hygroscopic).

Chemical Properties

1. Strong Acidic Nature: It is a strong diprotic (dibasic) acid, ionizing in two stages: $\text{H}_2\text{SO}_4 (\text{aq}) + \text{H}_2\text{O} (\text{l}) \rightleftharpoons \text{H}_3\text{O}^+ (\text{aq}) + \text{HSO}_4^- (\text{aq})$ $\text{HSO}_4^- (\text{aq}) + \text{H}_2\text{O} (\text{l}) \rightleftharpoons \text{H}_3\text{O}^+ (\text{aq}) + \text{SO}_4^{2-} (\text{aq})$ It reacts with bases, carbonates, and bicarbonates: $\text{2NaOH} (\text{aq}) + \text{H}_2\text{SO}_4 (\text{aq}) \rightarrow \text{Na}_2\text{SO}_4 (\text{aq}) + \text{2H}_2\text{O} (\text{l})$ $\text{Na}_2\text{CO}_3 (\text{s}) + \text{H}_2\text{SO}_4 (\text{dil}) \rightarrow \text{Na}_2\text{SO}_4 (\text{aq}) + \text{H}_2\text{O} (\text{l}) + \text{CO}_2 (\text{g})$

2. Oxidizing Agent (Concentrated $\text{H}_2\text{SO}_4$): Hot, concentrated sulfuric acid is a powerful oxidizing agent.

   • With non-metals: Oxidizes carbon to $\text{CO}_2$ and sulfur to $\text{SO}_2$. $\text{C} (\text{s}) + \text{2H}_2\text{SO}_4 (\text{conc}) \rightarrow \text{CO}_2 (\text{g}) + \text{2SO}_2 (\text{g}) + \text{2H}_2\text{O} (\text{l})$ $\text{S} (\text{s}) + \text{2H}_2\text{SO}_4 (\text{conc}) \rightarrow \text{3SO}_2 (\text{g}) + \text{2H}_2\text{O} (\text{l})$
   • With metals: Oxidizes moderately active metals (e.g., Cu, Zn) to their sulfates. $\text{Cu} (\text{s}) + \text{2H}_2\text{SO}_4 (\text{conc}) \rightarrow \text{CuSO}_4 (\text{aq}) + \text{SO}_2 (\text{g}) + \text{2H}_2\text{O} (\text{l})$ _Note: Cold, dilute $\text{H}_2\text{SO}4$ reacts with active metals (e.g., Zn, Mg, Fe) to produce hydrogen gas. $\text{Zn} (\text{s}) + \text{H}_2\text{SO}_4 (\text{dil}) \rightarrow \text{ZnSO}_4 (\text{aq}) + \text{H}_2 (\text{g})$

3. Dehydrating Agent: Concentrated sulfuric acid has a strong affinity for water and can remove water from various compounds.

   • From sugar: “Charring of sugar” $\text{C}{12}\text{H}{22}\text{O}_{11} (\text{s}) \xrightarrow{\text{conc. H}_2\text{SO}_4} \text{12C} (\text{s}) + \text{11H}_2\text{O} (\text{l})$
   • From formic acid: $\text{HCOOH} (\text{l}) \xrightarrow{\text{conc. H}_2\text{SO}_4} \text{CO} (\text{g}) + \text{H}_2\text{O} (\text{l})$
   • From oxalic acid: $(\text{COOH})_2 (\text{s}) \xrightarrow{\text{conc. H}_2\text{SO}_4} \text{CO} (\text{g}) + \text{CO}_2 (\text{g}) + \text{H}_2\text{O} (\text{l})$

Hydrogen Sulfide ($\text{H}_2\text{S}$)

Chemical Formula

$\text{H}_2\text{S}$

Common Name

Hydrogen Sulfide, Hydrosulfuric acid (in aqueous solution)

Laboratory Preparation

Hydrogen sulfide is typically prepared in the laboratory by the action of dilute mineral acids on a metal sulfide, commonly iron(II) sulfide.

Balanced Equation: $\text{FeS} (\text{s}) + \text{H}_2\text{SO}_4 (\text{dil}) \rightarrow \text{FeSO}_4 (\text{aq}) + \text{H}_2\text{S} (\text{g})$ Or $\text{FeS} (\text{s}) + \text{2HCl} (\text{dil}) \rightarrow \text{FeCl}_2 (\text{aq}) + \text{H}_2\text{S} (\text{g})$

Properties

Physical Properties

• Colorless gas with a characteristic rotten egg smell.
• Highly poisonous.
• Slightly soluble in water.

Chemical Properties

1. Weak Acidic Nature: Aqueous solution of $\text{H}_2\text{S}$ acts as a weak dibasic acid. $\text{H}_2\text{S} (\text{aq}) \rightleftharpoons \text{H}^+ (\text{aq}) + \text{HS}^- (\text{aq})$ $\text{HS}^- (\text{aq}) \rightleftharpoons \text{H}^+ (\text{aq}) + \text{S}^{2-} (\text{aq})$ It reacts with bases to form sulfides and bisulfides: $\text{NaOH} (\text{aq}) + \text{H}_2\text{S} (\text{g}) \rightarrow \text{NaHS} (\text{aq}) + \text{H}_2\text{O} (\text{l})$ (with excess $\text{H}_2\text{S}$) $\text{2NaOH} (\text{aq}) + \text{H}_2\text{S} (\text{g}) \rightarrow \text{Na}_2\text{S} (\text{aq}) + \text{2H}_2\text{O} (\text{l})$ (with excess $\text{NaOH}$)

2. Strong Reducing Agent: Hydrogen sulfide is a strong reducing agent because the sulfur atom is in its lowest oxidation state (-2) and can be easily oxidized to elemental sulfur (0) or higher oxidation states.

   • With sulfur dioxide: $\text{2H}_2\text{S} (\text{g}) + \text{SO}_2 (\text{g}) \rightarrow \text{3S} (\text{s}) + \text{2H}_2\text{O} (\text{l})$
   • With halogens: Reduces halogens to halide ions. $\text{H}_2\text{S} (\text{g}) + \text{Cl}_2 (\text{g}) \rightarrow \text{2HCl} (\text{g}) + \text{S} (\text{s})$
   • With oxidizing agents: Reduces acidified potassium permanganate (purple) to colorless manganese(II) ions, and acidified potassium dichromate (orange) to green chromium(III) ions. $\text{2KMnO}_4 + \text{5H}_2\text{S} + \text{3H}_2\text{SO}_4 \rightarrow \text{K}_2\text{SO}_4 + \text{2MnSO}_4 + \text{5S} + \text{8H}_2\text{O}$ $\text{K}_2\text{Cr}_2\text{O}_7 + \text{3H}_2\text{S} + \text{4H}_2\text{SO}_4 \rightarrow \text{K}_2\text{SO}_4 + \text{Cr}_2(\text{SO}_4)_3 + \text{3S} + \text{7H}_2\text{O}$

3. Precipitation of Metal Sulfides: This property is extensively used in qualitative analysis for the identification and separation of metal cations. Metal sulfides often have characteristic colors.

   • In acidic medium (Group II cations): $\text{CuSO}_4 (\text{aq}) + \text{H}_2\text{S} (\text{g}) \rightarrow \text{CuS} (\text{s}) \downarrow (\text{black}) + \text{H}_2\text{SO}_4 (\text{aq})$
   • In basic medium (Group IV cations): $\text{ZnSO}_4 (\text{aq}) + \text{H}_2\text{S} (\text{g}) \xrightarrow{\text{NH}_4\text{OH}} \text{ZnS} (\text{s}) \downarrow (\text{white}) + (\text{NH}_4)_2\text{SO}_4 (\text{aq})$

Comparative Properties of Sulfur Compounds

Property	Sulfur Dioxide ($\text{SO}_2$)	Sulfuric Acid ($\text{H}_2\text{SO}_4$)	Hydrogen Sulfide ($\text{H}_2\text{S}$)
Formula	$\text{SO}_2$	$\text{H}_2\text{SO}_4$	$\text{H}_2\text{S}$
Oxidation State of Sulfur	+4	+6	-2
Physical State (at STP)	Gas	Liquid	Gas
Acidity/Basicity	Acidic oxide	Strong diprotic acid	Weak dibasic acid
Redox Nature	Both oxidizing and reducing	Strong oxidizing (conc.), non-redox (dil.)	Strong reducing
Taste/Smell	Pungent, suffocating	Odorless (pure), corrosive	Rotten egg smell