Calculate the standard enthalpy change for this reaction using the following data.

Calculate the standard entropy change for this reaction using the following data.

The standard free energy change, Δ_G_θ, for the above reaction is –103 kJ mol–1 at 298 K.

Suggest why Δ_G_θ has a large negative value considering the sign of Δ_H_θ in part (a).

Deduce the ratio of Fe2+:Fe3+ in Fe3O4.

State the type of spectroscopy that could be used to determine their relative abundances.

b(i).

State the number of protons, neutrons and electrons in each species.

b(ii).

Iron has a relatively small specific heat capacity; the temperature of a 50 g sample rises by 44.4°C when it absorbs 1 kJ of heat energy.

Determine the specific heat capacity of iron, in J g−1 K−1. Use section 1 of the data booklet.

A voltaic cell is set up between the Fe2+ (aq) | Fe (s) and Fe3+ (aq) | Fe2+ (aq) half-cells.

Deduce the equation and the cell potential of the spontaneous reaction. Use section 24 of the data booklet.

The figure shows an apparatus that could be used to electroplate iron with zinc. Label the figure with the required substances.

Outline why, unlike typical transition metals, zinc compounds are not coloured.

Transition metals like iron can form complex ions. Discuss the bonding between transition metals and their ligands in terms of acid-base theory.

Estimate the time at which the powdered zinc was placed in the beaker.

a(i).

State what point Y on the graph represents.

a(ii).

The maximum temperature used to calculate the enthalpy of reaction was chosen at a point on the extrapolated (dotted) line.

State the maximum temperature which should be used and outline one assumption made in choosing this temperature on the extrapolated line.

Maximum temperature:

Assumption:

b(i).

To determine the enthalpy of reaction the experiment was carried out five times. The same volume and concentration of copper(II) sulfate was used but the mass of zinc was different each time. Suggest, with a reason, if zinc or copper(II) sulfate should be in excess for each trial.

b(ii).

The formula q = mcΔT was used to calculate the energy released. The values used in the calculation were m = 25.00 g, c = 4.18 J g−1 K−1.

State an assumption made when using these values for m and c.

b(iii).

Predict, giving a reason, how the final enthalpy of reaction calculated from this experiment would compare with the theoretical value.

b(iv).

Calculate the enthalpy change, in kJ, for the spray reaction, using the data below.

C 6 H 4 (OH) 2 (aq) → C 6 H 4 O 2 (aq) + H 2 (g) Δ H θ = + 177.0 kJ 2 H 2 O(l) + O 2 (g) → 2 H 2 O 2 (aq) Δ H θ = + 189.2 kJ H 2 O(l) → H 2 (g) + 1 2 O 2 (g) Δ H θ = + 285.5 kJ

The energy released by the reaction of one mole of hydrogen peroxide withhydroquinone is used to heat 850 cm3 of water initially at 21.8°C. Determine thehighest temperature reached by the water.

Specific heat capacity of water = 4.18 kJ kg−1 K−1.

(If you did not obtain an answer to part (i), use a value of 200.0 kJ for the energyreleased, although this is not the correct answer.)

Identify the species responsible for the peak at m/z = 110 in the mass spectrumof hydroquinone.

Identify the highest m/z value in the mass spectrum of quinone.

The diagram shows the Maxwell-Boltzmann curve for the uncatalyzed reaction.

Draw a distribution curve at a lower temperature (T2) and show on the diagram how the addition of a catalyst enables the reaction to take place more rapidly than at T1.

The hydrogen peroxide could cause further oxidation of the methanol. Suggest a possible oxidation product.

Determine the overall equation for the production of methanol.

c(i).

8.00 g of methane is completely converted to methanol. Calculate, to three significant figures, the final volume of hydrogen at STP, in dm3. Use sections 2 and 6 of the data booklet.

c(ii).

Determine the enthalpy change, Δ_H_, in kJ. Use section 11 of the data booklet.

Bond enthalpy of CO = 1077 kJ mol−1.

d(i).

State the expression for _K_c for this stage of the reaction.

d(ii).

State and explain the effect of increasing temperature on the value of Kc.

d(iii).

State the electron configuration of the Cu+ ion.

a(i).

Copper(II) chloride is used as a catalyst in the production of chlorine from hydrogen chloride.

4HCl (g) + O2 (g) → 2Cl2 (g) + 2H2O (g)

Calculate the standard enthalpy change, Δ_H_θ, in kJ, for this reaction, using section 12 of the data booklet.

a(ii).

The diagram shows the Maxwell–Boltzmann distribution and potential energy profile for the reaction without a catalyst.

Annotate both charts to show the activation energy for the catalysed reaction, using the label _E_a (cat).

a(iii).

Explain how the catalyst increases the rate of the reaction.

a(iv).

Solid copper(II) chloride absorbs moisture from the atmosphere to form a hydrate of formula CuCl2•xH2O.

A student heated a sample of hydrated copper(II) chloride, in order to determine the value of x. The following results were obtained:

Mass of crucible = 16.221 g
Initial mass of crucible and hydrated copper(II) chloride = 18.360 g
Final mass of crucible and anhydrous copper(II) chloride = 17.917 g

Determine the value of x.

State how current is conducted through the wires and through the electrolyte.

Wires:

Electrolyte:

c(i).

Write the half-equation for the formation of gas bubbles at electrode 1.

c(ii).

Bubbles of gas were also observed at another electrode. Identify the electrode and the gas.

Electrode number (on diagram):

Name of gas:

c(iii).

Deduce the half-equation for the formation of the gas identified in (c)(iii).

c(iv).

Determine the enthalpy of solution of copper(II) chloride, using data from sections 18 and 20 of the data booklet.

The enthalpy of hydration of the copper(II) ion is −2161 kJ mol−1.

Calculate the cell potential at 298 K for the disproportionation reaction, in V, using section 24 of the data booklet.

e(i).

Comment on the spontaneity of the disproportionation reaction at 298 K.

e(ii).

Calculate the standard Gibbs free energy change, Δ_G_θ, to two significant figures, for the disproportionation at 298 K. Use your answer from (e)(i) and sections 1 and 2 of the data booklet.

e(iii).

Suggest, giving a reason, whether the entropy of the system increases or decreases during the disproportionation.

e(iv).

Deduce, giving a reason, the sign of the standard enthalpy change, Δ_H_θ, for the disproportionation reaction at 298 K.

e(v).

Predict, giving a reason, the effect of increasing temperature on the stability of copper(I) chloride solution.

e(vi).

Describe how the blue colour is produced in the Cu(II) solution. Refer to section 17 of the data booklet.

f(i).

Deduce why the Cu(I) solution is colourless.

f(ii).

When excess ammonia is added to copper(II) chloride solution, the dark blue complex ion, [Cu(NH3)4(H2O)2]2+, forms.

State the molecular geometry of this complex ion, and the bond angles within it.

Molecular geometry:

Bond angles:

f(iii).

Examine the relationship between the Brønsted–Lowry and Lewis definitions of a base, referring to the ligands in the complex ion [CuCl4]2−.

f(iv).