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Chemistry · Ch 7 — Redox Reactions

Decomposition reactions

2

Decomposition reactions

Decomposition reactions are the reverse of combination reactions. A compound breaks down into two or more simpler substances, and at least one of the products must be in the elemental state (oxidation number zero).

Examples:

2H2+1O−2(l)→Δ2H20(g)+O20(g)(7.26) 2\overset{+1}{H_2}\overset{-2}{O}(l) \xrightarrow{\Delta} 2\overset{0}{H_2}(g) + \overset{0}{O_2}(g) \qquad(7.26)

Hydrogen is reduced (+1 → 0), oxygen is oxidised (–2 → 0).

2Na+1H−1(s)→Δ2Na0(s)+H20(g)(7.27) 2\overset{+1}{Na}\overset{-1}{H}(s) \xrightarrow{\Delta} 2\overset{0}{Na}(s) + \overset{0}{H_2}(g) \qquad(7.27)

Sodium is reduced (+1 → 0), hydrogen is oxidised (–1 → 0).

2K+1Cl+5O3−2(s)→Δ2K+1Cl−1(s)+3O20(g)(7.28) 2\overset{+1}{K}\overset{+5}{Cl}\overset{-2}{O_3}(s) \xrightarrow{\Delta} 2\overset{+1}{K}\overset{-1}{Cl}(s) + 3\overset{0}{O_2}(g) \qquad(7.28)

Chlorine is reduced (+5 → –1), oxygen is oxidised (–2 → 0). Potassium's oxidation number stays at +1 throughout — no change.

Important

Not all decomposition reactions are redox reactions. The key test: does at least one product contain an element in its elemental form? If not, it is not a redox reaction.

Example of a non-redox decomposition: …

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