[Maximum number: 2]

Copper is a transition element and has atomic number 29.

(a)

The following equilibrium exists between two complex ions of copper in the +2 oxidation state.

\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+4 \mathrm{Cl}^{-} \rightleftharpoons\left[\mathrm{CuCl}_{4}\right]^{2-}+6 \mathrm{H}_{2} \mathrm{O}

[ 1 ]

(i)

Write the expression for the stability constant, $K_{\text {stab }}$ , for this equilibrium.

K_{\text {stab }}=

[ 1 ]

(b)

Copper also forms the complex ions $\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2+}$ and $\left[\mathrm{Cu}(e n)\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2+}$ where en is the bidentate ligand ethane-1,2-diamine, $\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}$ .

\begin{gathered} {\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+2 \mathrm{NH}_{3} \rightleftharpoons\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2+}+2 \mathrm{H}_{2} \mathrm{O}} \\ {\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+e n \rightleftharpoons\left[\mathrm{Cu}(e n)\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2+}+2 \mathrm{H}_{2} \mathrm{O}} \end{gathered}

equilibrium 1
equilibrium 2

[ 1 ]

(i)

The table lists the values of stability constants for these two complexes.

What do these $K_{\text {stab }}$ values tell us about the relative positions of equilibria 1 and 2?

[ 1 ]

[Maximum number: 3]

$1 \mathrm{EDTA}^{4-}$ , is a polydentate ligand.

(a)

(i)

Write an expression for the stability constant, $K_{\text {stab1 }}$ , for equilibrium 1, and state its units.

K_{\text {stab1 }}=

units =

[ 2 ]

(b)

Cadmium ions form complexes with methylamine, $\mathrm{CH}_{3} \mathrm{NH}_{2}$ , and with 1,2-diaminoethane, $\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}$ , as shown in equilibriums 2 and 3. 1,2-diaminoethane is shown as en.
equilibrium $2\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+4 \mathrm{CH}_{3} \mathrm{NH}_{2} \rightleftharpoons\left[\mathrm{Cd}\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{H}_{2} \mathrm{O} \quad K_{\text {stab2 }}=3.60 \times 10^{6}$
equilibrium $3\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+2 \mathrm{en} \rightleftharpoons\left[\mathrm{Cd}(\mathrm{en})_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{H}_{2} \mathrm{O} \quad \mathrm{K}_{\text {stab3 }}=4.20 \times 10^{10}$

An equilibrium is set up between these two complexes as shown in equilibrium 4.
equilibrium $4\left[\mathrm{Cd}\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+2 \mathrm{en} \rightleftharpoons\left[\mathrm{Cd}(\mathrm{en})_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{CH}_{3} \mathrm{NH}_{2}$

\begin{aligned} & \Delta H^{\ominus}=+0.840 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ & \Delta S^{\ominus}=+80.9 \mathrm{JK}^{-1} \mathrm{~mol}^{-1} \end{aligned}

[ 1 ]

(i)

$K_{\text {eq } 4}$ is the equilibrium constant for equilibrium 4.

Write an expression for $K_{\text {eq4 }}$ in terms of $K_{\text {stab2 }}$ and $K_{\text {stab3 }}$ .
$K_{\text {eq } 4}=$

[ 1 ]

(a)

Consider the following two equilibria and associated data values at 298 K .

\begin{array}{ccc} \mathrm{AgBr}(\mathrm{~s}) \rightleftharpoons \mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{Br}^{-}(\mathrm{aq}) & \text { equilibrium } 1 \quad K_{\mathrm{sp}}=5.0 \times 10^{-13} \mathrm{~mol}^{2} \mathrm{dm}^{-6} \\ \mathrm{Ag}^{+}(\mathrm{aq})+2 \mathrm{NH}_{3}(\mathrm{aq}) \rightleftharpoons\left[\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}\right]^{+}(\mathrm{aq}) & \text { equilibrium } 2 \quad K_{\text {stab }}=1.7 \times 10^{7} \mathrm{~mol}^{-2} \mathrm{dm}^{6} \end{array}

The equilibrium constant for equilibrium 1 is the solubility product, $K_{\mathrm{sp}}$ , of AgBr(s). The equilibrium constant for equilibrium 2 is the stability constant, $K_{\text {stab }}$ , for the formation of $\left[\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}\right]^{+}(\mathrm{aq})$ .

(i)

Use equilibria 1 and 2 to construct an equation for the reaction of AgBr(s) with concentrated $\mathrm{NH}_{3}(\mathrm{aq})$ . This is equilibrium 3.

(ii)

Write an expression for the equilibrium constant of equilibrium 3, $K_{\text {eq3 }}$ , in terms of $K_{\mathrm{sp}}$ for equilibrium 1 and $K_{\text {stab }}$ for equilibrium 2.

K_{\text {eq3 }}=

(a)

(i)

Define stability constant, $K_{\text {stab }}$ .

[ 1 ]

(ii)

Nickel can form complexes with the ligands en, $\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}$ , and $t n, \mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}$ , as shown.
equilibrium $1\left[\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+3 e n \rightleftharpoons\left[\mathrm{Ni}(e n)_{3}\right]^{2+}+6 \mathrm{H}_{2} \mathrm{O} \quad K_{\text {stab }}=6.76 \times 10^{17}$
equilibrium $2 \quad\left[\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+3 t n \rightleftharpoons\left[\mathrm{Ni}(t n)_{3}\right]^{2+}+6 \mathrm{H}_{2} \mathrm{O} \quad K_{\text {stab }}=1.86 \times 10^{12}$
Construct an expression for the stability constant, $K_{\text {stab }}$ , for equilibrium 1. State the units for $K_{\text {stab }}$ .
$K_{\text {stab }}=$

units =

[ 2 ]

(iii)

Describe what the $K_{\text {stab }}$ values indicate about the position of equilibrium for equilibrium 1 and 2. Use the $K_{\text {stab }}$ values to deduce which complex, $\left[\mathrm{Ni}(e n)_{3}\right]^{2+}$ or $\left[\mathrm{Ni}(t n)_{3}\right]^{2+}$ , is more stable.

[ 1 ]

(a)

EDTA ${ }^{4-}$ is a polydentate ligand.

[ 2 ]

(i)

The complex [CaEDTA] ${ }^{2-}$ can be used to remove toxic metals from the body.

Table 1.2 shows the numerical values for the stability constants, $K_{\text {stab }}$ , for some metal ions with EDTA ${ }^{4-}$ .

Table 1.2

An aqueous solution containing [CaEDTA] ${ }^{2-}$ is added to a solution containing equal concentrations of $\mathrm{Cr}^{3+}(\mathrm{aq}), \mathrm{Fe}^{3+}(\mathrm{aq})$ and $\mathrm{Pb}^{2+}(\mathrm{aq})$ . The resulting mixture is left to reach a state of equilibrium.

State the type of reaction when $[\mathrm{CaEDTA}]^{2-}$ reacts with $\mathrm{Cr}^{3+}(\mathrm{aq}), \mathrm{Fe}^{3+}(\mathrm{aq})$ and $\mathrm{Pb}^{2+}(\mathrm{aq})$ .

[ 1 ]

(ii)

Deduce the relative concentrations of $[\mathrm{CrEDTA}]^{-},[\mathrm{FeEDTA}]^{-}$ and $[\mathrm{PbEDTA}]^{2-}$ present in the resulting mixture.

Explain your answer. > >
highest concentration
lowest concentration

[ 1 ]

(a)

(i)

State what is meant by the term stability constant.

[ 1 ]

(ii)

Complete the table by placing one tick ( ✓ ) in each row to suggest how increasing temperature will affect $K_{\text {stab }}$ and the equilibrium concentration of the cadmium complex, $\left[\left[\mathrm{Cd}\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}\right]$ , for equilibrium 1. Explain your answer.

explanation

EDTA ${ }^{4-}$ is a polydentate ligand. When a solution of EDTA ${ }^{4-}$ is added to $\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}$ a new complex [CdEDTA] ${ }^{2-}$ is formed.

The values for the stability constants for two $\mathrm{Cd}^{2+}$ complexes are shown.

[ 2 ]

(iii)

A solution containing equal numbers of moles of $\mathrm{CH}_{3} \mathrm{NH}_{2}$ and EDTA is added to $\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}$ . Predict which complex is formed in the larger amount. Explain your answer.

[ 1 ]

(a)

Copper can form complexes with the ligands ammonia and en, $\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}$ , as shown. $\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}(\mathrm{aq})+\mathrm{en}(\mathrm{aq}) \rightleftharpoons\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}(\mathrm{en})\right]^{2+}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{I}) \quad K_{\text {stab }}=3.98 \times 10^{10}$ equilibrium 4 $\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}(\mathrm{aq})+2 \mathrm{NH}_{3}(\mathrm{aq}) \rightleftharpoons\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\left(\mathrm{NH}_{3}\right)_{2}\right]^{2+}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{I}) \quad K_{\text {stab }}=5.01 \times 10^{7}$ equilibrium 5

[ 3 ]

(i)

Write an expression for the stability constant, $K_{\text {stab }}$ , for equilibrium 5. State its units.

K_{\text {stab }}=

units =

[ 2 ]

(ii)

Of the three copper complexes in equilibria 4 and 5, state the formula of the copper complex that is the most stable and explain your choice.
copper complex explanation

[ 1 ]

(a)

The numerical values for the stability constants, $K_{\text {stab }}$ , of two other silver(I) complexes are given.

An aqueous solution containing $\mathrm{Ag}^{+}$ is added to a solution containing equal concentrations of $\mathrm{CN}^{-}(\mathrm{aq}), \mathrm{NH}_{3}(\mathrm{aq})$ and $\mathrm{S}_{2} \mathrm{O}_{3}{ }^{2-}(\mathrm{aq})$ . The mixture is left to reach equilibrium.

Deduce the relative concentrations of $\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-},\left[\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}\right]^{+}$ and $\left[\mathrm{Ag}\left(\mathrm{S}_{2} \mathrm{O}_{3}\right)_{2}\right]^{3-}$ present in the resulting mixture. Explain your answer. > >
highest concentration
lowest concentration

[ 2 ]

(a)

Transition metal ions often exist as hexa-aqua complexes in aqueous solution. The reactions which involve ligand exchange are reversible.

\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+4 \mathrm{NH}_{3} \rightleftharpoons\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{H}_{2} \mathrm{O}

[ 3 ]

(i)

Write an expression for the stability constant, $K_{\text {stab }}$ , for this equilibrium. Give its units.

K_{\text {stab }}=

units =

[ 2 ]

(ii)

The numerical value for $K_{\text {stab }}$ for this equilibrium at 298 K is $1.20 \times 10^{13}$ .

Explain how this value relates to the relative stabilities of the two complexes.

[ 1 ]

[Maximum number: 6]

Bubbling air through different aqueous mixtures of $\mathrm{CoCl}_{2}, \mathrm{NH}_{4} \mathrm{Cl}$ and $\mathrm{NH}_{3}$ produces various complex ions with the general formula $\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{6-\mathrm{n}} \mathrm{Cl}_{\mathrm{n}}\right]^{3-\mathrm{n}}$ .

(a)

Iron(III) forms complexes in separate reactions with both $\mathrm{SCN}^{-}$ ions and $\mathrm{Cl}^{-}$ ions.

\begin{aligned} \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SCN}^{-}(\mathrm{aq}) & \rightleftharpoons\left[\mathrm{FeSCN}^{2+}(\mathrm{aq})\right. \\ \mathrm{Fe}^{3+}(\mathrm{aq})+4 \mathrm{Cl}^{-}(\mathrm{aq}) & \rightleftharpoons\left[\mathrm{FeCl}_{4}\right]^{-}(\mathrm{aq}) \end{aligned}

[ 6 ]

(i)

Write the expressions for the stability constants, $K_{\text {stab }}$ , for these two equilibria. Include units in your answers.

K_{\text {stab1 }}=

units =

K_{\text {stab2 }}=

units =

[ 3 ]

(ii)

An equilibrium can be set up between these two complexes as shown in equilibrium 3 .

\left[\mathrm{FeCl}_{4}\right]^{-}(\mathrm{aq})+\mathrm{SCN}^{-}(\mathrm{aq}) \rightleftharpoons[\mathrm{FeSCN}]^{2+}(\mathrm{aq})+4 \mathrm{Cl}^{-}(\mathrm{aq})

Write an expression for $K_{\text {eq3 }}$ in terms of $K_{\text {stab1 }}$ and $K_{\text {stab2 }}$ .
$K_{\text {eq } 3}=$

[ 1 ]

(iii)

The numerical values for these stability constants are shown.

K_{\text {stab1 }}=1.4 \times 10^{2} \quad K_{\text {stab2 }}=8.0 \times 10^{-2}

Calculate the value of $K_{\text {eq } 3}$ stating its units.
$K_{\text {eq } 3}=$ units =

[ 2 ]