Define stability constant in equilibrium terms.
Calculate the oxidation states of Cu in $[\text{Cu(EDTA)}]^{2-}$ and Cr in $[\text{Cr(EDTA)}]^-$.
Deduce how many lone pairs are donated by each $\text{EDTA}^{4-}$ ligand in one $[\text{Fe(EDTA)}]^{2-}$ complex ion.
Identify the most stable complex in Table 4.1. Explain why you selected it.
At equilibrium in a solution at $298\,\text{K}$, $[[\text{Cu(H}_2\text{O)}_6]^{2+}] = 3.00 \times 10^{-10}\,\text{mol dm}^{-3}$ and $[\text{EDTA}^{4-}] = 5.00 \times 10^{-12}\,\text{mol dm}^{-3}$. Use the expression for $K_{\text{stab}}$ to determine the concentration of $[\text{Cu(EDTA)}]^{2-}$ in this solution. Show your working.
A solution containing $[\text{Cu(EDTA)}]^{2-}$ ions is pale blue, whereas a solution containing $[\text{Cu(NH}_3)_4(\text{H}_2\text{O})_2]^{2+}$ ions is deep blue. Explain this difference in colour.