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HL Paper 2

This question is about the decomposition of hydrogen peroxide.

Hydrogen peroxide decomposes to water and oxygen when a catalyst such as potassium iodide, KI, is added.

2H₂O₂ (aq) $\xrightarrow{{{\text{KI (aq)}}}}$ O₂ (g) + 2H₂O (l)

Suggest why many chemicals, including hydrogen peroxide, are kept in brown bottles instead of clear colourless bottles.

[1]

In a laboratory experiment solutions of potassium iodide and hydrogen peroxide were mixed and the volume of oxygen generated was recorded. The volume was adjusted to 0 at t = 0.

The data for the first trial is given below.

Plot a graph on the axes below and from it determine the average rate of
formation of oxygen gas in cm³ O₂ (g) s⁻¹.

Average rate of reaction:

[3]

b(i).

Two more trials (2 and 3) were carried out. The results are given below.

Determine the rate equation for the reaction and its overall order, using your answer from (b)(i).

Rate equation:

Overall order:

[2]

b(ii).

Additional experiments were carried out at an elevated temperature. On the axes below, sketch Maxwell–Boltzmann energy distribution curves at two temperatures T₁ and T₂, where T₂ > T₁.

[2]

b(iii).

Apart from a greater frequency of collisions, explain, by annotating your graphs in (b)(iii), why an increased temperature causes the rate of reaction to increase.

[2]

b(iv).

MnO₂ is another possible catalyst for the reaction. State the IUPAC name for MnO₂.

[1]

b(v).

Comment on why peracetic acid, CH3COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) ⇌ CH₃COOOH (aq) + H₂O (l)

[1]

Sodium percarbonate, 2Na₂CO₃•3H₂O₂, is an adduct of sodium carbonate and hydrogen peroxide and is used as a cleaning agent.

M_r (2Na₂CO₃•3H₂O₂) = 314.04

Calculate the percentage by mass of hydrogen peroxide in sodium percarbonate, giving your answer to two decimal places.

[2]

Rhenium, Re, was the last element with a stable isotope to be isolated.

Before its isolation, scientists predicted the existence of rhenium and some of its properties.

One chloride of rhenium has the empirical formula ReCl₃.

Rhenium forms salts containing the perrhenate(VII) ion, ReO₄⁻.

The stable isotope of rhenium contains 110 neutrons.

State the nuclear symbol notation ${}_{\text{Z}}^{\text{A}}{\text{X}}$ for this isotope.

[1]

Suggest the basis of these predictions.

[2]

b(i).

A scientist wants to investigate the catalytic properties of a thin layer of rhenium metal on a graphite surface.

Describe an electrochemical process to produce a layer of rhenium on graphite.

[2]

b(ii).

Predict two other chemical properties you would expect rhenium to have, given its position in the periodic table.

[2]

b(iii).

Describe how the relative reactivity of rhenium, compared to silver, zinc, and copper, can be established using pieces of rhenium and solutions of these metal sulfates.

[2]

State the name of this compound, applying IUPAC rules.

[1]

d(i).

Calculate the percentage, by mass, of rhenium in ReCl₃.

[2]

d(ii).

Suggest why the existence of salts containing an ion with this formula could be predicted. Refer to section 6 of the data booklet.

[1]

e(i).

Deduce the coefficients required to complete the half-equation.

ReO₄⁻ (aq) + ____H⁺ (aq) + ____e⁻ ⇌ [Re(OH)₂]²⁺ (aq) + ____H₂O (l) E^θ = +0.36 V

[1]

e(ii).

Predict, giving a reason, whether the reduction of ReO₄⁻ to [Re(OH)₂]²⁺ would oxidize Fe²⁺ to Fe³⁺ in aqueous solution. Use section 24 of the data booklet.

[1]

e(iii).

When heated in air, magnesium ribbon reacts with oxygen to form magnesium oxide.

The reaction in (a)(i) was carried out in a crucible with a lid and the following data was recorded:

Mass of crucible and lid = 47.372 ±0.001 g

Mass of crucible, lid and magnesium ribbon before heating = 53.726 ±0.001 g

Mass of crucible, lid and product after heating = 56.941 ±0.001 g

When magnesium is burnt in air, some of it reacts with nitrogen to form magnesium nitride according to the equation:

3 Mg (s) + N₂(g) → Mg₃N₂(s)

The presence of magnesium nitride can be demonstrated by adding water to the product. It is hydrolysed to form magnesium hydroxide and ammonia.

Most nitride ions are ¹⁴N^3–.

Write a balanced equation for the reaction that occurs.

[1]

a(i).

Identify a metal, in the same period as magnesium, that does not form a basic oxide.

[1]

a(ii).

Calculate the amount of magnesium, in mol, that was used.

[1]

b(i).

Determine the percentage uncertainty of the mass of product after heating.

[2]

b(ii).

Assume the reaction in (a)(i) is the only one occurring and it goes to completion, but some product has been lost from the crucible. Deduce the percentage yield of magnesium oxide in the crucible.

[2]

b(iii).

Evaluate whether this, rather than the loss of product, could explain the yield found in (b)(iii).

[1]

c(i).

Suggest an explanation, other than product being lost from the crucible or reacting with nitrogen, that could explain the yield found in (b)(iii).

[1]

c(ii).

Calculate coefficients that balance the equation for the following reaction.

[1]

d(i).

Ammonia is added to water that contains a few drops of an indicator. Identify an indicator that would change colour. Use sections 21 and 22 of the data booklet.

[1]

d(ii).

Determine the oxidation state of nitrogen in Mg₃N₂ and in NH₃.

[1]

d(iii).

Deduce, giving reasons, whether the reaction of magnesium nitride with water is an acid–base reaction, a redox reaction, neither or both.

[2]

d(iv).

State the number of subatomic particles in this ion.

[1]

e(i).

Some nitride ions are ¹⁵N^3–. State the term that describes the relationship between ¹⁴N^3– and ¹⁵N^3–.

[1]

e(ii).

The nitride ion and the magnesium ion are isoelectronic (they have the same electron configuration). Determine, giving a reason, which has the greater ionic radius.

[1]

e(iii).

Suggest, giving a reason, whether magnesium or nitrogen would have the greater sixth ionization energy.

[1]

e(iv).

Suggest two reasons why atoms are no longer regarded as the indivisible units of matter.

[2]

State the types of bonding in magnesium, oxygen and magnesium oxide, and how the valence electrons produce these types of bonding.

[4]

Copper forms two chlorides, copper(I) chloride and copper(II) chloride.

Two electrolysis cells were assembled using graphite electrodes and connected in series as shown.

Copper(I) chloride undergoes a disproportionation reaction, producing copper(II) chloride and copper.

2Cu⁺ (aq) → Cu (s) + Cu²⁺ (aq)

Dilute copper(II) chloride solution is light blue, while copper(I) chloride solution is colourless.

State the electron configuration of the Cu⁺ ion.

[1]

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, ΔH^θ, in kJ, for this reaction, using section 12 of the data booklet.

[2]

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)}.

[2]

a(iii).

Explain how the catalyst increases the rate of the reaction.

[2]

a(iv).

Solid copper(II) chloride absorbs moisture from the atmosphere to form a hydrate of formula CuCl₂• $x$ H₂O.

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 $x$ .

[3]

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

Wires:

Electrolyte:

[2]

c(i).

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

[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:

[1]

c(iii).

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

[1]

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.

[2]

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

[1]

e(i).

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

[1]

e(ii).

Calculate the standard Gibbs free energy change, ΔG^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.

[1]

e(iii).

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

[1]

e(iv).

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

[1]

e(v).

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

[1]

e(vi).

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

[3]

f(i).

Deduce why the Cu(I) solution is colourless.

[1]

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:

[1]

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⁻.

[2]

f(iv).