Energy Cycles in Reactions (DP IB Chemistry: HL): Exam Questions

Nitrogen oxides produced by combustion are largely nitrogen monoxide or nitrogen dioxide.

i) Draw Lewis diagrams for nitrogen monoxide and nitrogen dioxide.

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ii) Using the diagrams, explain the meaning of the term free radical.

Platinum and rhodium are found in catalytic converters and facilitate the conversion of Carbon monoxide and nitrogen monoxides to nitrogen and carbon dioxide.

i) Write an equation for this reaction.

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ii) State the changes in oxidation state for each carbon and nitrogen.

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Using your answer to part (bi) and the average bond enthalpies from section 12 of the data booklet,determine the enthalpy change for the reaction between carbon monoxide and nitrogen monoxide.

Hess's Law provides a method for calculating enthalpy changes for reactions that cannot be measured directly by constructing an alternative reaction pathway. This often involves the use of standard enthalpy of formation (ΔH_f^ө) data.

State Hess’s Law.

Define the term standard enthalpy of formation, ΔH^ϴ_f.

The following equation represents a step in the extraction of titanium:

TiCl₄ (g) + 4Na (l) → 4NaCl (s) + Ti (s)

The enthalpy change for this reaction can be calculated using the standard enthalpy of formation (ΔH^ө_f) data shown in Table 1.

Table 1 

Calculate the standard enthalpy change for this reaction, ΔH^ϴ_r.

For some reactions, it is more useful to visualize the alternative pathway.

Construct a Hess's Law cycle using enthalpies of formation for the reaction of calcium fluoride with sulfuric acid.

CaF₂ (s) + H₂SO₄ (aq) → 2HF (g) + CaSO₄ (s)

Define the term standard enthalpy of combustion, ΔH^ϴ_c.

Write an equation for the complete combustion of propan-1-ol, CH₃CH₂CH₂OH (l).

Construct a Hess’s Law cycle for the complete combustion of propan-1-ol.

The enthalpy of formation of propan-1-ol is -303 kJ mol^-1.

Using section 13 of the data booklet, calculate the enthalpy change of the reaction, ΔH^ϴ_r.

The industrial production of hydrogen cyanide can be achieved by the reaction of methane and ammonia.

i) Write the balanced chemical equation for the conversion of one mole of methane and one mole of ammonia to form hydrogen cyanide and hydrogen.

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ii) Draw the Lewis structure for the hydrogen cyanide molecule, HCN.

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Use Hess’s Law and the information below to calculate the enthalpy change for this reaction.

Using your answer to part (b), draw a labelled reaction profile diagram for this reaction.

Define the term standard enthalpy of reaction, ΔH^ϴ_r.

Butane, C₄H₁₀,  is typically used as fuel for cigarette lighters and portable stoves, a propellant in aerosols, a heating fuel, a refrigerant, and in the manufacture of a wide range of products.

Write a balanced chemical equation for the complete combustion of butane.

The enthalpy change for the hydrogenation of butene to form butane is shown below.

C₄H₈ + H₂  → C₄H₁₀

Using section 12 of the data booklet, calculate the enthalpy of hydrogenation of butene.

The standard enthalpy of formation of butane, ΔH^ө_f(C₄H₁₀), can be determined using a Hess's Law cycle and enthalpy of combustion data.

i) Write the chemical equation that represents the standard enthalpy of formation of butane.

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ii) Construct a Hess's Law cycle to relate the enthalpy of formation of butane to the enthalpies of combustion of carbon, hydrogen, and butane.

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iii) Using the data below, calculate the standard enthalpy of formation of butane.

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The accepted value for the enthalpy of hydrogenation of butene is -30.3 kJ mol^-1.

Suggest one reason why your value calculated using average bond enthalpies in part (b) is different from this accepted value.

The enthalpy of combustion for ethene gas can be determined theoretically using average bond enthalpies.

Write the balanced chemical equation for the complete combustion of ethene, C₂H₄, assuming liquid water is formed.

Define the term average bond enthalpy.

Using your answer to (a) and section 12 of the data booklet, calculate the enthalpy of combustion of ethene.

Bond enthalpies can be found using Hess’s Law or from experimental data.

Outline the difference between the two ways of finding bond enthalpy.

Pure crystals of lithium fluoride are used in X-ray monochromator.

i) Define the term enthalpy of atomisation.

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ii) Explain why the enthalpy of atomisation of fluorine is positive.

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iii) Complete the Born–Haber cycle for lithium fluoride by adding the missing species on the lines.

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Use the data in the following table and your completed Born–Haber cycle from part (a) to answer the questions below.

i) Calculate the enthalpy of lattice formation of lithium fluoride.

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ii) Explain and justify how the enthalpy of lattice formation of LiBr compares with that of LiF. You must refer to the size of the ions in your answer.

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Calcium chloride has many uses including as an agent to lower the freezing point of water. It is very effective for preventing ice formation on road surfaces and as a deicer.

i) Define the term ionisation energy.

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ii) Explain why the second ionisation energy of calcium is greater than the first ionisation energy.

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Describe the structure and bonding in calcium chloride.

The Born-Haber cycle for CaCl₂ is shown:

Using section 9 in the data booklet and the following information, calculate the enthalpy change for the following conversions. 

• ΔH^θ_IE2 Ca = 1145 kJ mol^-1                       

• ΔH^θ_at Ca = 178 kJ mol^-1

• ΔH^θ_BE Cl₂ = 242 kJ mol^-1

i) Ca (s) → Ca²⁺ (g) + 2e^-

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ii) Cl₂ (g) + 2e^- → 2Cl^- (g)

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Using section 16 of the data booklet, calculate the value for the enthalpy of formation for calcium chloride, ΔH^θ_f CaCl₂. 

A Born-Haber cycle can be used to determine the electron affinity of fluorine using data for magnesium fluoride, MgF₂.

Define the term electron affinity in relation to fluorine.

The table below shows the enthalpy data required.  

Suggest why the first electron affinity of fluorine is an exothermic change.

Vanadium is commonly found in different ores such as magnetite, vanadinite and patronite. The vanadium is commonly extracted from these ores by reduction and displacement.

Vanadium can be extracted by the reduction of vanadium pentoxide, V₂O₅, with calcium at high temperatures, according to the following equation.

V₂O₅ (s) + 5Ca (s) → 2V (s) + 5CaO (s)

The enthalpy of formation of vanadium pentoxide is -1560 kJ mol^-1 and the standard enthalpy change for the reaction is -1615 kJ mol^-1. 

Construct a Hess’s Law cycle for this reaction.

Use the data in part a) to calculate the enthalpy of formation, ΔH_f, of calcium oxide in kJ mol^-1. 

Explain why calcium is able to reduce vanadium(V) oxide in this reaction.

The compound diborane, B₂H₆, is used as a rocket fuel. The equation for the combustion of diborane is shown below.

B₂H₆ (g) + 3O₂ (g) → B₂O₃ (s) + 3H₂O (l)

Calculate the standard enthalpy change of this reaction using the following data

I. 2B (s) + 3H₂ (g) → B₂H₆ (g)          ΔH = 36 kJ mol^-1

II. H₂ (g) + ½O₂ (g) → H₂O (l)          ΔH = -286 kJ mol^-1

III. 2B (s) + 1½O₂ (g) → B₂O₃ (s)       ΔH = -1274 kJ mol^-1

Ethyne, C₂H₂, is a highly combustible gas used in welding.

Write the equation for the complete combustion of ethyne gas.

The standard enthalpies of formation for the substances involved are provided in section 13 of the data booklet.

i) Construct a Hess's Law cycle to determine the standard enthalpy of combustion of ethyne, ΔH_c, using enthalpy of formation data.

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ii) Using your cycle and the data booklet, calculate the standard enthalpy of combustion of ethyne.

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Coal can be converted into "coal gas", a fuel containing carbon monoxide and hydrogen.

The overall equation for coal gasification is

H₂O (l) + C (s) → CO (g) + H₂ (g) 

Using the equations below, calculate the standard enthalpy change for this reaction. 

I. 2C (s) + O₂ (g) → 2CO (g)            ΔH = -222 kJ

II. 2H₂ (g) + O₂ (g) → 2H₂O (g)        ΔH = -484 kJ

III. H₂O (l) → H₂O (g)                       ΔH = +44 kJ

The coal gas produced is then combusted:

CO (g) + H₂ (g) + O₂ (g) → CO₂ (g) + H₂O (g) 

Calculate the standard enthalpy change for this combustion, using the following data.

I. 2C (s) + O₂ (g) → 2CO (g)       ΔH = -222 kJ

II. C (s) + O₂ (g) → CO₂ (g)         ΔH = -394 kJ

III. 2H₂ (g) + O₂ (g) → 2H₂O (g) ΔH = -484 kJ

E85 is a biofuel blend containing a high percentage of ethanol in gasoline.

If E85 is vaporised before combustion, predict if the energy released would be greater than, less than, or the same as combusting liquid E85. Explain your reasoning.

A 1.00 kg sample of E85 is found to contain 60.0% ethanol by mass. Using sections 7 and 14 of the data booklet:

i) Calculate the amount, in moles, of ethanol in the sample.

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ii) Calculate the energy released, in kJ, from the complete combustion of the ethanol in the sample.

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Strontium salts have a number of applications such as fireworks, flares, glow in the dark paint and toothpaste for sensitive teeth. The strontium required for these salts can be extracted from the ore strontia, SrO, by displacement with powdered aluminium in a vacuum.

i) Write a balanced symbol equation, including state symbols, for the reaction of strontia with aluminium.

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ii) State the role of the aluminium in this reaction.

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The standard enthalpy change for this extraction of strontium is 99.3 kJ mol^-1 and the standard enthalpy of formation of aluminium oxide is -1676.7 kJ mol^-1

Calculate the standard enthalpy of formation, ΔH_f, in kJ mol^-1 of strontia.

Manganese is too brittle for use as a pure metal, so it is often alloyed with other metals. Manganese is used in steel to increase the strength and resistance to wear. Manganese steel (13% Mn) is extremely strong and used for railway tracks, safes and prison bars. Alloys of 1.5% manganese with aluminium are used to make drinks cans due to the improved corrosion resistance of the alloy.

Manganese is extracted from different ores by reduction with carbon monoxide.

Mn₂O₃ (s) + 3CO (g) → 2Mn (s) + 3CO₂ (g)

The enthalpy of formation, ΔH_f, of Mn₂O₃ (s) is −971 kJ mol^-1. Use this information and section 13 of the data booklet to calculate the enthalpy change of reaction, ΔH_r, in kJ mol^-1. 

The reaction in part c) reaches equilibrium at high temperatures.

Using your answer to part c, explain how temperature can be altered to increase the yield of the reaction and explain the effect that this would have on the rate of reaction.

Fluorine, the most electronegative element, forms a compound with oxygen called oxygen difluoride, OF₂.

Define the term average bond enthalpy.

The standard enthalpy change for the formation of oxygen difluoride is +28 kJ mol^-1.

F₂ (g) + ½O₂ (g) → OF₂ (g)

Using section 12 of the data booklet, determine the bond enthalpy of the O–F bond.

Ethanol reacts with ethanoyl chloride, CH₃COCl, to produce the ester ethyl ethanoate and hydrogen chloride gas.

Write the balanced chemical equation for this reaction.

Using section 12 of the data booklet, calculate the overall enthalpy change for this reaction.

Methane reacts violently with fluorine to form carbon tetrafluoride and hydrogen fluoride

Write the equation for this reaction. 

Using your answer to part a) and section 12 of the data booklet, calculate the standard enthalpy change for this reaction

Sketch a labelled energy diagram for the reaction of methane and fluorine.

Lattice enthalpies can be determined experimentally using a Born–Haber cycle and theoretically using calculations based on electrostatic principles.

The experimental lattice enthalpies of magnesium chloride, MgCl₂ , calcium chloride, CaCl₂ , strontium chloride, SrCl₂ , and barium chloride, BaCl₂ are given in section 16 of the data booklet. Explain the trend in the values. 

Explain why strontium chloride, SrCl₂ , has a much greater lattice enthalpy than rubidium chloride, RbCl.

Strontium is used as a red colouring agent in fireworks as it provides a very intense red colour. Use sections 9 and 16 to calculate the enthalpy of atomisation for chlorine in strontium chloride.

The incomplete Born-Haber cycle for sodium selenide is shown below. 

State the equations for processes 1, 2 and 3. 

Using the data provided and values from section 9 of the data booklet, calculate the lattice enthalpy of formation for sodium selenide. [3]

If sulfur is used as opposed to selenium in the lattice, what would you expect to happen to the value of the enthalpy of lattice dissociation. Explain your answer.