HL Paper 2
Chlorine undergoes many reactions.
of manganese(IV) oxide was added to of .
Chlorine gas reacts with water to produce hypochlorous acid and hydrochloric acid.
is a common chlorofluorocarbon, .
State the full electron configuration of the chlorine atom.
State, giving a reason, whether the chlorine atom or the chloride ion has a larger radius.
Outline why the chlorine atom has a smaller atomic radius than the sulfur atom.
The mass spectrum of chlorine is shown.
NIST Mass Spectrometry Data Center Collection © 2014 copyright by the U.S. Secretary of Commerce on behalf of the United States of America. All rights reserved.
Outline the reason for the two peaks at and .
Explain the presence and relative abundance of the peak at .
Calculate the amount, in , of manganese(IV) oxide added.
Determine the limiting reactant, showing your calculations.
Determine the excess amount, in , of the other reactant.
Calculate the volume of chlorine, in , produced if the reaction is conducted at standard temperature and pressure (STP). Use section 2 of the data booklet.
State the oxidation state of manganese in and .
Deduce, referring to oxidation states, whether is an oxidizing or reducing agent.
Hypochlorous acid is considered a weak acid. Outline what is meant by the term weak acid.
State the formula of the conjugate base of hypochlorous acid.
Calculate the concentration of in a solution with a .
State the type of reaction occurring when ethane reacts with chlorine to produce chloroethane.
Predict, giving a reason, whether ethane or chloroethane is more reactive.
Explain the mechanism of the reaction between chloroethane and aqueous sodium hydroxide, , using curly arrows to represent the movement of electron pairs.
Ethoxyethane (diethyl ether) can be used as a solvent for this conversion.
Draw the structural formula of ethoxyethane
Deduce the number of signals and chemical shifts with splitting patterns in the 1H NMR spectrum of ethoxyethane. Use section 27 of the data booklet.
Calculate the percentage by mass of chlorine in .
Comment on how international cooperation has contributed to the lowering of emissions responsible for ozone depletion.
s produce chlorine radicals. Write two successive propagation steps to show how chlorine radicals catalyse the depletion of ozone.
Calcium carbide, CaC2, is an ionic solid.
Describe the nature of ionic bonding.
Describe how the relative atomic mass of a sample of calcium could be determined from its mass spectrum.
When calcium compounds are introduced into a gas flame a red colour is seen; sodium compounds give a yellow flame. Outline the source of the colours and why they are different.
Suggest two reasons why solid calcium has a greater density than solid potassium.
Outline why solid calcium is a good conductor of electricity.
Sketch a graph of the first six ionization energies of calcium.
Calcium carbide reacts with water to form ethyne and calcium hydroxide.
CaC2(s) + H2O(l) → C2H2(g) + Ca(OH)2(aq)
Estimate the pH of the resultant solution.
Describe how sigma (σ) and pi () bonds are formed.
Deduce the number of σ and bonds in a molecule of ethyne.
Bromine can form the bromate(V) ion, BrO3−.
State the electron configuration of a bromine atom.
Sketch the orbital diagram of the valence shell of a bromine atom (ground state) on the energy axis provided. Use boxes to represent orbitals and arrows to represent electrons.
Draw two Lewis (electron dot) structures for BrO3−.
Determine the preferred Lewis structure based on the formal charge on the bromine atom, giving your reasons.
Predict, using the VSEPR theory, the geometry of the BrO3− ion and the O−Br−O bond angles.
Bromate(V) ions act as oxidizing agents in acidic conditions to form bromide ions.
Deduce the half-equation for this reduction reaction.
Bromate(V) ions oxidize iron(II) ions, Fe2+, to iron(III) ions, Fe3+.
Deduce the equation for this redox reaction.
Calculate the standard Gibbs free energy change, ΔGΘ, in J, of the redox reaction in (ii), using sections 1 and 24 of the data booklet.
EΘ (BrO3− / Br−) = +1.44 V
State and explain the magnetic property of iron(II) and iron(III) ions.
Magnesium is a group 2 metal which exists as a number of isotopes and forms many compounds.
Magnesium ions produce no emission or absorption lines in the visible region of the electromagnetic spectrum. Suggest why most magnesium compounds tested in a school laboratory show traces of yellow in the flame.
(i) Explain the convergence of lines in a hydrogen emission spectrum.
(ii) State what can be determined from the frequency of the convergence limit.
Magnesium chloride can be electrolysed.
(i) Deduce the half-equations for the reactions at each electrode when molten magnesium chloride is electrolysed, showing the state symbols of the products. The melting points of magnesium and magnesium chloride are 922K and 987K respectively.
(ii) Identify the type of reaction occurring at the cathode (negative electrode).
(iii) State the products when a very dilute aqueous solution of magnesium chloride is electrolysed.
Standard electrode potentials are measured relative to the standard hydrogen electrode. Describe a standard hydrogen electrode.
A magnesium half-cell, Mg(s)/Mg2+(aq), can be connected to a copper half-cell, Cu(s)/Cu2+(aq).
(i) Formulate an equation for the spontaneous reaction that occurs when the circuit is completed.
(ii) Determine the standard cell potential, in V, for the cell. Refer to section 24 of the data booklet.
(iii) Predict, giving a reason, the change in cell potential when the concentration of copper ions increases.
Nitric acid is usually produced by the oxidation of ammonia.
A mixture of nitric acid and sulfuric acid can be used to convert benzene to nitrobenzene, C6H5NO2.
Draw arrows in the boxes to represent the electron configuration of a nitrogen atom.
Deduce a Lewis (electron dot) structure of the nitric acid molecule, HNO3, that obeys the octet rule, showing any non-zero formal charges on the atoms.
Explain the relative lengths of the three bonds between N and O in nitric acid.
State a technique used to determine the length of the bonds between N and O in solid HNO3.
Write an equation for the reaction between the acids to produce the electrophile, NO2+.
Draw the structural formula of the carbocation intermediate produced when this electrophile attacks benzene.
Deduce the number of signals that you would expect in the 1H NMR spectrum of nitrobenzene and the relative areas of these.
This question is about iron.
Deduce the full electron configuration of Fe2+.
Explain why, when ligands bond to the iron ion causing the d-orbitals to split, the complex is coloured.
State the nuclear symbol notation, , for iron-54.
Mass spectrometry analysis of a sample of iron gave the following results:
Calculate the relative atomic mass, Ar, of this sample of iron to two decimal places.
An iron nail and a copper nail are inserted into a lemon.
Explain why a potential is detected when the nails are connected through a voltmeter.
Calculate the standard electrode potential, in V, when the Fe2+ (aq) | Fe (s) and Cu2+ (aq) | Cu (s) standard half-cells are connected at 298 K. Use section 24 of the data booklet.
Calculate ΔGθ, in kJ, for the spontaneous reaction in (f)(i), using sections 1 and 2 of the data booklet.
Calculate a value for the equilibrium constant, Kc, at 298 K, giving your answer to two significant figures. Use your answer to (f)(ii) and section 1 of the data booklet.
(If you did not obtain an answer to (f)(ii), use −140 kJ mol−1, but this is not the correct value.)
Properties of elements and their compounds can be related to the position of the elements in the periodic table.
Explain the decrease in atomic radius from Na to Cl.
Explain why the radius of the sodium ion, Na+, is smaller than the radius of the oxide ion, O2−.
Sketch a graph to show the relative values of the successive ionization energies of boron.
Predict, giving your reasons, whether Mn2+ or Fe2+ is likely to have a more exothermic enthalpy of hydration.
Fast moving helium nuclei (4He2+) were fired at a thin piece of gold foil with most passing undeflected but a few deviating largely from their path. The diagram illustrates this historic experiment.
Figure from PPLATO / FLAP (Flexible Learning Approach To Physics), http://www.met.reading.ac.uk/pplato2/h-flap/
phys8_1.html#top 1996 The Open University and The University of Reading.
Suggest what can be concluded about the gold atom from this experiment.
Subsequent experiments showed electrons existing in energy levels occupying various orbital shapes.
Sketch diagrams of 1s, 2s and 2p.
State the electron configuration of copper.
Copper is a transition metal that forms different coloured complexes. A complex [Cu(H2O)6]2+ (aq) changes colour when excess Cl− (aq) is added.
Explain the cause of this colour change, using sections 3 and 15 from the data booklet.
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 14N3–.
Write a balanced equation for the reaction that occurs.
Identify a metal, in the same period as magnesium, that does not form a basic oxide.
Calculate the amount of magnesium, in mol, that was used.
Determine the percentage uncertainty of the mass of product after heating.
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.
Evaluate whether this, rather than the loss of product, could explain the yield found in (b)(iii).
Suggest an explanation, other than product being lost from the crucible or reacting with nitrogen, that could explain the yield found in (b)(iii).
Calculate coefficients that balance the equation for the following reaction.
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.
Determine the oxidation state of nitrogen in Mg3N2 and in NH3.
Deduce, giving reasons, whether the reaction of magnesium nitride with water is an acid–base reaction, a redox reaction, neither or both.
State the number of subatomic particles in this ion.
Some nitride ions are 15N3–. State the term that describes the relationship between 14N3– and 15N3–.
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.
Suggest, giving a reason, whether magnesium or nitrogen would have the greater sixth ionization energy.
Suggest two reasons why atoms are no longer regarded as the indivisible units of matter.
State the types of bonding in magnesium, oxygen and magnesium oxide, and how the valence electrons produce these types of bonding.
Iron may be extracted from iron (II) sulfide, FeS.
Iron (II) sulfide, FeS, is ionically bonded.
The first step in the extraction of iron from iron (II) sulfide is to roast it in air to form iron (III) oxide and sulfur dioxide.
Outline why metals, like iron, can conduct electricity.
Justify why sulfur is classified as a non-metal by giving two of its chemical properties.
Sketch the first eight successive ionisation energies of sulfur.
Describe the bonding in this type of solid.
State a technique that could be used to determine the crystal structure of the solid compound.
State the full electron configuration of the sulfide ion.
Outline, in terms of their electronic structures, why the ionic radius of the sulfide ion is greater than that of the oxide ion.
Suggest why chemists find it convenient to classify bonding into ionic, covalent and metallic.
Write the equation for this reaction.
Deduce the change in the oxidation state of sulfur.
Suggest why this process might raise environmental concerns.
Explain why the addition of small amounts of carbon to iron makes the metal harder.
Magnetite, Fe3O4, is another ore of iron that contains both Fe2+ and Fe3+.
Iron exists as several isotopes.
Deduce the ratio of Fe2+:Fe3+ in Fe3O4.
State the type of spectroscopy that could be used to determine their relative abundances.
State the number of protons, neutrons and electrons in each species.
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.
Oxygen exists as two allotropes, diatomic oxygen, O2, and ozone, O3.
Draw a Lewis (electron dot) structure for ozone.
Discuss the relative length of the two O−O bonds in ozone.
Explain why there are frequencies of UV light that will dissociate O3 but not O2.
Explain, using equations, how the presence of results in a chain reaction that decreases the concentration of ozone in the stratosphere.
Analytical chemistry uses instruments to separate, identify, and quantify matter.
Nitric oxide reacts with chlorine.
2NO (g) + Cl2 (g) → 2NOCl (g)
The following experimental data were obtained at 101.3 kPa and 263 K.
Menthol is an organic compound containing carbon, hydrogen and oxygen.
Outline how this spectrum is related to the energy levels in the hydrogen atom.
A sample of magnesium has the following isotopic composition.
Calculate the relative atomic mass of magnesium based on this data, giving your answer to two decimal places.
Complete combustion of 0.1595 g of menthol produces 0.4490 g of carbon dioxide and 0.1840 g of water. Determine the empirical formula of the compound showing your working.
0.150 g sample of menthol, when vaporized, had a volume of 0.0337 dm3 at 150 °C and 100.2 kPa. Calculate its molar mass showing your working.
Determine the molecular formula of menthol using your answers from parts (d)(i) and (ii).
Deduce the order of reaction with respect to Cl2 and NO.
State the rate expression for the reaction.
Calculate the value of the rate constant at 263 K.
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.
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, ΔHθ, in kJ, for this reaction, using section 12 of the data booklet.
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 Ea (cat).
Explain how the catalyst increases the rate of the reaction.
Solid copper(II) chloride absorbs moisture from the atmosphere to form a hydrate of formula CuCl2•H2O.
A student heated a sample of hydrated copper(II) chloride, in order to determine the value of . 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 .
State how current is conducted through the wires and through the electrolyte.
Wires:
Electrolyte:
Write the half-equation for the formation of gas bubbles at electrode 1.
Bubbles of gas were also observed at another electrode. Identify the electrode and the gas.
Electrode number (on diagram):
Name of gas:
Deduce the half-equation for the formation of the gas identified in (c)(iii).
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.
Comment on the spontaneity of the disproportionation reaction at 298 K.
Calculate the standard Gibbs free energy change, ΔGθ, 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.
Suggest, giving a reason, whether the entropy of the system increases or decreases during the disproportionation.
Deduce, giving a reason, the sign of the standard enthalpy change, ΔHθ, for the disproportionation reaction at 298 K.
Predict, giving a reason, the effect of increasing temperature on the stability of copper(I) chloride solution.
Describe how the blue colour is produced in the Cu(II) solution. Refer to section 17 of the data booklet.
Deduce why the Cu(I) solution is colourless.
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:
Examine the relationship between the Brønsted–Lowry and Lewis definitions of a base, referring to the ligands in the complex ion [CuCl4]2−.
Dinitrogen monoxide, N2O, causes depletion of ozone in the stratosphere.
Different sources of N2O have different ratios of 14N : 15N.
The Lewis (electron dot) structure of the dinitrogen monoxide molecule can be represented as:
Outline why ozone in the stratosphere is important.
Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).
Write two equations to show how NO (g) catalyses the decomposition of ozone.
State one analytical technique that could be used to determine the ratio of 14N : 15N.
A sample of gas was enriched to contain 2 % by mass of 15N with the remainder being 14N.
Calculate the relative molecular mass of the resulting N2O.
Predict, giving two reasons, how the first ionization energy of 15N compares with that of 14N.
Explain why the first ionization energy of nitrogen is greater than both carbon and oxygen.
Nitrogen and carbon:
Nitrogen and oxygen:
State what the presence of alternative Lewis structures shows about the nature of the bonding in the molecule.
State, giving a reason, the shape of the dinitrogen monoxide molecule.
Deduce the hybridization of the central nitrogen atom in the molecule.
Iron(II) disulfide, FeS2, has been mistaken for gold.
State the full electronic configuration of Fe2+.
Explain why there is a large increase from the 8th to the 9th ionization energy of iron.
Calculate the oxidation state of sulfur in iron(II) disulfide, FeS2.
Describe the bonding in iron, Fe (s).
The emission spectrum of an element can be used to identify it.
Hydrogen spectral data give the frequency of 3.28 × 1015 s−1 for its convergence limit.
Calculate the ionization energy, in J, for a single atom of hydrogen using sections 1 and 2 of the data booklet.
Calculate the wavelength, in m, for the electron transition corresponding to the frequency in (a)(iii) using section 1 of the data booklet.
Deduce any change in the colour of the electrolyte during electrolysis.
Deduce the gas formed at the anode (positive electrode) when graphite is used in place of copper.
Explain why transition metals exhibit variable oxidation states in contrast to alkali metals.
The properties of elements can be predicted from their position in the periodic table.
Explain why Si has a smaller atomic radius than Al.
Explain why the first ionization energy of sulfur is lower than that of phosphorus.
State the condensed electron configurations for Cr and Cr3+.
Describe metallic bonding and how it contributes to electrical conductivity.
Deduce, giving a reason, which complex ion [Cr(CN)6]3− or [Cr(OH)6]3− absorbs higher energy light. Use section 15 of the data booklet.
[Cr(OH)6]3− forms a green solution. Estimate a wavelength of light absorbed by this complex, using section 17 of the data booklet.
Deduce the Lewis (electron dot) structure and molecular geometry of sulfur tetrafluoride, SF4, and sulfur dichloride, SCl2.
Suggest, giving reasons, the relative volatilities of SCl2 and H2O.
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 for this isotope.
Suggest the basis of these predictions.
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.
Predict two other chemical properties you would expect rhenium to have, given its position in the periodic table.
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.
State the name of this compound, applying IUPAC rules.
Calculate the percentage, by mass, of rhenium in ReCl3.
Suggest why the existence of salts containing an ion with this formula could be predicted. Refer to section 6 of the data booklet.
Deduce the coefficients required to complete the half-equation.
ReO4− (aq) + ____H+ (aq) + ____e− ⇌ [Re(OH)2]2+ (aq) + ____H2O (l) Eθ = +0.36 V
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.