Sulfur (S): Key Reactions & Properties

Chemical Properties Overview

Sulfur (S) is a non-metal belonging to Group 16 (Chalcogens) of the periodic table, located in Period 3.

Atomic Number: 16
Electronic Configuration: $[Ne] 3s^2 3p^4$
Electronegativity (Pauling): 2.58
Common Oxidation States: -2, 0, +2, +4, +6
Reactivity: Generally forms covalent compounds. It is less electronegative than oxygen and less reactive than fluorine and oxygen but more reactive than halogens in certain contexts (e.g., forming H₂S). It readily combines with metals and non-metals.

Sulfur burns in air or oxygen with a blue flame, producing sulfur dioxide.

Combustion in Oxygen: $$S_{(s)} + O_{2(g)} \xrightarrow{\text{heat}} SO_{2(g)}$$
Catalytic Oxidation of Sulfur Dioxide (Contact Process intermediate): Sulfur dioxide can be further oxidized to sulfur trioxide, especially in the presence of catalysts like vanadium(V) oxide ($V_2O_5$) and high temperature, a key step in the Contact Process for sulfuric acid manufacturing. $$2SO_{2(g)} + O_{2(g)} \xrightarrow{V_2O_5, \text{ heat}} 2SO_{3(g)}$$

Elemental sulfur is generally insoluble in water and does not react with water or steam under normal conditions to form new compounds.

Sulfur is unreactive with non-oxidizing dilute acids. However, it reacts with strong oxidizing acids.

With Concentrated Hot Sulfuric Acid ($H_2SO_4$): Sulfur is oxidized to sulfur dioxide, and sulfuric acid is reduced. $$S_{(s)} + 2H_2SO_{4(conc)} \xrightarrow{\text{hot}} 3SO_{2(g)} + 2H_2O_{(l)}$$
With Concentrated Hot Nitric Acid ($HNO_3$): Sulfur is oxidized to sulfuric acid, and nitric acid is reduced to nitrogen dioxide. $$S_{(s)} + 6HNO_{3(conc)} \xrightarrow{\text{hot}} H_2SO_{4(aq)} + 6NO_{2(g)} + 2H_2O_{(l)}$$

Sulfur undergoes disproportionation (simultaneous oxidation and reduction) when heated with concentrated alkalis.

With Hot Concentrated Sodium Hydroxide ($NaOH$): Sulfur is simultaneously oxidized to sulfite ($SO_3^{2-}$) and reduced to sulfide ($S^{2-}$). $$3S_{(s)} + 6NaOH_{(aq)} \xrightarrow{\text{hot}} 2Na_2S_{(aq)} + Na_2SO_{3(aq)} + 3H_2O_{(l)}$$

Lead Acetate Test: When treated with lead acetate solution, sulfide ions produce a black precipitate of lead(II) sulfide. $$Pb^{2+}{(aq)} + S^{2-}{(aq)} \rightarrow PbS_{(s)} \text{ (Black precipitate)}$$ (Alternatively, a filter paper moistened with lead acetate solution turns black when exposed to $H_2S$ gas).
Sodium Nitroprusside Test: Sulfide ions give a violet coloration with freshly prepared sodium nitroprusside solution. $$Na_2S_{(aq)} + Na_2[Fe(CN)5NO]{(aq)} \rightarrow Na_4[Fe(CN)5NOS]{(aq)} \text{ (Violet coloration)}$$
Dilute Acid Test: On adding dilute HCl or H₂SO₄, sulfide salts evolve hydrogen sulfide ($H_2S$) gas, which has a characteristic rotten egg smell. $$S^{2-}{(aq)} + 2H^+{(aq)} \rightarrow H_2S_{(g)}$$

Dilute Acid Test: On adding dilute HCl or H₂SO₄, sulfite salts evolve sulfur dioxide ($SO_2$) gas, which has a suffocating smell. $SO_2$ turns acidified potassium dichromate solution green (due to the formation of $Cr^{3+}$) and turns lime water milky. $$SO_3^{2-}{(aq)} + 2H^+{(aq)} \rightarrow SO_{2(g)} + H_2O_{(l)}$$ $$K_2Cr_2O_{7(aq)} + H_2SO_{4(aq)} + 3SO_{2(g)} \rightarrow K_2SO_{4(aq)} + Cr_2(SO_4){3(aq)} \text{ (Green)} + H_2O{(l)}$$
Barium Chloride Test: Sulfite ions react with barium chloride solution to form a white precipitate of barium sulfite, which is soluble in dilute HCl. $$Ba^{2+}{(aq)} + SO_3^{2-}{(aq)} \rightarrow BaSO_{3(s)} \text{ (White precipitate, soluble in dilute HCl)}$$

Barium Chloride Test: Sulfate ions react with barium chloride solution to form a white precipitate of barium sulfate, which is insoluble in dilute HCl and dilute HNO₃. This test distinguishes sulfate from sulfite. $$Ba^{2+}{(aq)} + SO_4^{2-}{(aq)} \rightarrow BaSO_{4(s)} \text{ (White precipitate, insoluble in dilute HCl/HNO}_3\text{)}$$
Lead Acetate Test: Sulfate ions form a white precipitate of lead(II) sulfate with lead acetate solution. $$Pb^{2+}{(aq)} + SO_4^{2-}{(aq)} \rightarrow PbSO_{4(s)} \text{ (White precipitate)}$$