Give the meaning of standard electrode potential.
Three redox systems, A, B and C, are displayed. The ligand 1,2-diaminoethane, $\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2$, is shown as en. Two electrochemical cells are assembled to compare the standard electrode potentials, $E^\circ$, of three half-cells. The diagrams indicate the relative potential of each electrode. Use this information to fill in the table by inserting A, B and C and work out the order of $E^\circ$ for the three half-cells.
The complex $[\text{Ru(en)}_3]^{3+}$ exhibits stereoisomerism. The ligand en is bidentate. Draw three-dimensional diagrams to show the two isomers of $[\text{Ru(en)}_3]^{3+}$. State the type of stereoisomerism.
An electrochemical cell has a $\text{Br}_2/\text{Br}^-$ half-cell and an $\text{Ag}^+/\text{Ag}$ half-cell, both under standard conditions. Use the Data Booklet to calculate $E^\circ_{\text{cell}}$. Deduce the direction of electron flow in the wire passing through the voltmeter between these two half-cells.
Water is added to the $\text{Ag}^+/\text{Ag}$ half-cell in (b)(i). Predict the effect this has on $E_{\text{cell}}$. Put a tick (✓) in the correct box: less positive, no change, more positive. Explain your response.
Silver bromide, $\text{AgBr}$, dissolves in aqueous $\text{S}_2\text{O}_3^{2-}$ ions to produce the complex ion $[\text{Ag(S}_2\text{O}_3)_2]^{3-}$. The $\text{S}_2\text{O}_3^{2-}$ ions function as monodentate ligands. For equilibrium 1, $\text{AgBr(s)} + 2\text{S}_2\text{O}_3^{2-}\text{(aq)} \rightleftharpoons [\text{Ag(S}_2\text{O}_3)_2]^{3-}\text{(aq)} + \text{Br}^-\text{(aq)}$, define ligand.
Write the expression for the equilibrium constant, $K_c$, for equilibrium 1.
The solubility product, $K_{sp}$, of $\text{AgBr}$ is $5.4 \times 10^{-13}$ and the stability constant, $K_{stab}$, of $[\text{Ag(S}_2\text{O}_3)_2]^{3-}$ is $2.9 \times 10^{13}$. Use your result from part (c)(ii) together with these data to calculate $K_c$ for equilibrium 1. Give the units for $K_c$.
The stability constants, $K_{stab}$, for two other silver(I) complexes are provided numerically: $[\text{Ag(CN)}_2]^-$ has $5.3 \times 10^{18}$ and $[\text{Ag(NH}_3)_2]^+$ has $1.6 \times 10^7$. An aqueous $\text{Ag}^+$ solution is added to a solution containing equal concentrations of $\text{CN}^-\text{(aq)}$, $\text{NH}_3\text{(aq)}$ and $\text{S}_2\text{O}_3^{2-}\text{(aq)}$. Deduce the relative amounts of $[\text{Ag(CN)}_2]^-$, $[\text{Ag(NH}_3)_2]^+$ and $[\text{Ag(S}_2\text{O}_3)_2]^{3-}$ present in the final mixture. Explain your reasoning.