Energy Cycles in Reactions (HL IB Chemistry)

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

Draw Lewis diagrams for nitrogen monoxide and nitrogen dioxide and use the diagrams to 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.

Write an equation for the reaction and state the changes in oxidation state for each carbon and nitrogen.

Use your answer to part (c) and the bond enthalpy data given in Table 1 to determine the enthalpy change for the reaction between carbon monoxide and nitrogen monoxide.

Table 1

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

State Hess’s Law.

The following equation represents the second step in the extraction of titanium, using the Kroll process:

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

Use the standard formation data shown in Table 1 to calculate the enthalpy change for the reaction, ΔH^ϴ_r.

Table 1 

Construct a Hess’s Law cycle for the reaction of calcium fluoride, CaF₂ (s) , and sulfuric acid, H₂SO₄ (aq).

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 propanol, CH₃CH₂CH₂OH (l).

Construct a Hess’s Law cycle for the complete combustion of propanol.

Table 1

Use the data given in Table 1 in part (d) to calculate the enthalpy change of the reaction, ΔH^ϴ_r.

Urea can be used as a fertiliser and is manufactured by the reaction of ammonia and carbon dioxide via the following equation. 

2NH₃ (g) + CO₂ (g) → NH₂CONH₂ (s) + H₂O (l)

Using the data in Table 1 calculate the enthalpy change for the formation of urea, ΔH_r^ө.

Table 1

Ammonia reacts with oxygen to produce steam and nitrogen(II) oxide. Draw a Hess’s Law cycle which could be used to calculate the enthalpy change of the reaction using formation data.

Use Hess’s Law and the information below to calculate the enthalpy change, ΔH^ϴ_r, for the conversion of one mole of ethene and one mole of hydrogen to one mole of ethane. 

C₂H₄ (g) + 3O₂ (g) → 2CO₂ (g) + 2H₂O (l)                  ΔH^ϴ_r = -1411 kJ mol^-1

C₂H₆ (g) + 3.5O₂ (g) → 2CO₂ (g) + 3H₂O (l)               ΔH^ϴ_r = -1560 kJ mol^-1

H₂ (g) + 0.5O₂ (g) → H₂O (l)                                       ΔH^ϴ_r = –286 kJ mol^-1

Use Hess’s Law and the information below to calculate the enthalpy change for the conversion of one mole of solid carbon into carbon monoxide.

C (s) + O₂ (g) → CO₂  (g)                    ΔH^ϴ_r = - 393.5 kJ mol^-1

CO (g) + ½O₂ (g)  → CO₂  (g)            ΔH^ϴ_r = - 283.5 kJ mol^-1

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

Use Hess’s Law and the information below to calculate the enthalpy change, ΔH^ϴ_r, for the conversion of methane and ammonia to form hydrogen cyanide and hydrogen.

N₂ (g) + 3H₂ (g) → 2NH₃ (g)                           ΔH^ϴ_r = -91.8 kJ

C (s) + 2H₂ (g) → CH₄ (g)                                        ΔH^ϴ_r = -74.9 kJ

H₂ (g) + 2C (g) + N₂ (g) → 2HCN (g)              ΔH^ϴ_r = 270.3 kJ

Using your answer to part (b) draw a reaction profile diagram for the reaction outlined.

Draw the Lewis structure for hydrogen cyanide, HCN.

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 an equation for the complete combustion of butane.

Determine the enthalpy of formation of butane, C₄H₁₀ , using the enthalpy of combustion data below.

C (s) + O₂ (g) → CO₂ (g)                                            ΔH^ϴ_f = -394 kJ           

H₂ (g) + 0.5O₂ (g) → H₂O (l)                                       ΔH^ϴ_f = -286 kJ

C₄H₁₀ (g) + 6.5O₂ (g) → 4CO₂ (g) + 5H₂O (l)             ΔH^ϴ_f = -2878 kJ

Butane can be formed from the hydrogenation of butene. Using the data in Table 1, determine a value for the enthalpy of formation.

Table 1

Enthalpy changes can be found using bond enthalpy data. Some bond enthalpy values are shown below in Table 1.

Table 1

 The balanced equation for the reaction between methane and propane is

CH₄ (g) +  CH₃CH₂CH₃ (g) → CH₃CH₂CH₂CH₃ (g) + H₂ (g)

Use the equation and bond enthalpy data to calculate the enthalpy change for the above reaction.                                  

Define the term average bond enthalpy.

Enthalpy changes can be found using bond enthalpy data. Some bond enthalpy values are shown below in Table 2. 

Table 2

The balanced equation for the combustion of ethene is

C₂H₄ (g) + 3O₂ (g) → 2CO₂ (g) + 2H₂O (l)

Use the equation and bond enthalpy data to 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.

Use the energy level diagram to determine the activation energy, E_a, for the given reaction in Figure 1.

Figure 1

Ethene can be hydrated via the following reaction:

C₂H₄ (g) + H₂O (g) → C₂H₅OH (g)

Table 1

Use the data in Table 1 to calculate the enthalpy change for the hydration of ethene.

Explain why the value to your answer to part (b) is different from the data book value for the hydration of ethene.

C₂H₄ (g) + H₂O (g) → C₂H₅OH (g)

Table 2 below has some enthalpy data for a different chemical reaction. Hydrazine, N₂H₄ can react with hydrogen peroxide in an exothermic reaction, as shown below.

N₂H₄(g) + 2H₂O₂ → N₂(g) + 4H₂O(g)                          ΔH^ϴ_r = -789 kJ mol^-1

The structure of hydrazine is shown in Figure 1.

Figure 1

Table 2

Using the reaction equation and the data in the table above, calculate the value of the N-H bond in hydrazine.

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

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

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.

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

Using Section 16 of the Data Booklet, calculate the value for the enthalpy of formation for calcium chloride, ΔH^θ_f CaCl₂. 

This question is about fluorine and the associated energy changes when it reacts with magnesium to form magnesium fluoride.

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. 

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

1. 2B (s) + 3H₂ (g) → B₂H₆ (g)          ΔH = 36 kJ mol^-1
2. H₂ (g) + ½O₂ (g) → H₂O (l)          ΔH = -286 kJ mol^-1
3. 2B (s) + 1½O₂ (g) → B₂O₃ (s)       ΔH = -1274 kJ mol^-1

Ethyne, C₂H₂, is a useful gas as it gives a high temperature flame when burnt with oxygen. State the equation for the combustion of ethyne gas.

Use your answer to part b) to construct a Hess's Law cycle for the combustion of ethyne gas.

Use section 13 in the data booklet to determine the enthalpy of combustion, ΔH_c, of ethyne gas. 

Coal gasification converts coal into a combustible mixture of carbon monoxide and hydrogen known as coal gas, in a gasifier. 

H₂O (l) + C (s) → CO (g) + H₂ (g) 

Using the following equations, calculate the enthalpy change of reaction, ΔH_r, in kJ for cola gasification. 

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

This coal gas can be used as a fuel as the following equation shows.

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

Calculation the enthalpy change of reaction, ΔH_r, in kJ for this combustion reaction from the following equations.

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

Blending amounts of alternative fuel with conventional fuel is one way to replace petroleum. A fuel blend of 51% to 83% ethanol and the remaining being gasoline is known as E85.

If the fuel blend is vaporised before combustion, predict whether the amount of energy released would be greater, less or the same. Explain your answer.

Use sections 7 and 14 of the Data booklet to calculate the following.

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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.

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

Use this information to 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.

Use your answer to part c) to 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.

Define the term average bond enthalpy.

Determine the bond dissociation energy, in kJ mol^-1, for one mole of O−F bonds using the following equation and section 12 of the data booklet. Give your answer to 3 significant figures. 

F₂ (g) + ½O₂ (g) → OF₂ (g)    ΔH_r = +28 kJ mol^-1

The reaction of ethanoyl chloride, CH₃COCl , and ethanol form an ester. State the equation for this reaction.

Use section 12 in the data booklet to deduce the energy required, in kJ mol^-1, to break the bonds.

Deduce the energy released, in kJ mol^-1, when the bonds are formed and therefore the enthalpy change for the reaction.

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

Formulate the equation for this reaction. 

Use your answer to part a) and section 12 of the data booklet to calculate the following:

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

Hydrazine has the formula N₂H₄ and is used as a rocket fuel (e.g. for the Apollo moon rockets). It burns in the following reaction for which the enthalpy change is -583 kJ mol^-1.

N₂H₄ (g) + O₂ (g) → N₂ (g) + 2H₂O (g)

Sketch the Lewis structure of hydrazine, N₂H₄.

Use section 12 of the Data booklet and the information in part a) to deduce the bond enthalpy, in kJ mol^-1, for the N-N bond.

Outline why the value of enthalpy of reaction calculated from bond enthalpies is less accurate.

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. 

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.

Use section 9 in the data booklet and the information in the table to calculate the lattice enthalpy of aluminium oxide.