Iron can form stable ions in the $+2$ and $+3$ oxidation states. Explain why transition elements show variable oxidation states.
Iron(II) salts in aqueous solution contain the complex ion $[\mathrm{Fe(H_2O)_6}]^{2+}$. Define complex ion.
$[\mathrm{Fe(H_2O)_6}]^{2+}$ can be changed into $[\mathrm{Fe(H_2O)_4(OH)_2}]$. Suggest a suitable reagent for this change. State the type of reaction.
$[\mathrm{Fe(H_2O)_4(OH)_2}]$ is a green precipitate that becomes brown on standing in air. Table 2.1 lists electrode potentials for these half-equations: - $\mathrm{Fe(H_2O)_3(OH)_3 + H_2O + e^- \rightleftharpoons Fe(H_2O)_4(OH)_2 + OH^-}$, $E^\circ = -0.56\ \mathrm{V}$ - $\mathrm{O_2 + 2H_2O + 4e^- \rightleftharpoons 4OH^-}$, $E^\circ = +0.40\ \mathrm{V}$ Use the information in Table 2.1 to explain why $[\mathrm{Fe(H_2O)_4(OH)_2}]$ turns brown on standing in air. Include an equation for this reaction.
The complex $[\mathrm{Co(NH_3)_6}]^{2+}$ undergoes reaction with hydrogen peroxide as shown. Reaction 1: $2[\mathrm{Co(NH_3)_6}]^{2+} + \mathrm{H_2O_2} \rightarrow 2[\mathrm{Co(NH_3)_6}]^{3+} + 2\mathrm{OH^-}$ $E^\circ_{\mathrm{cell}} = +1.67\ \mathrm{V}$ Calculate $\Delta G^\circ$, in $\mathrm{kJ\ mol^{-1}}$, for reaction 1.