Energy Cycles (DP IB Chemistry: HL)

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

i)

Define the term enthalpy of atomisation

ii)

Explain why the enthalpy of atomisation of fluorine is positive

iii)

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

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

This question is about enthalpy changes in solution.

Explain the decrease why the value for the enthalpy of hydration, ΔH^θ_hyd, of group 1 ions increases from lithium to caesium. 

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 8 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 18 of the Data Booklet, calculate the value for the enthalpy of formation for calcium chloride, ΔH^θ_f CaCl₂. 

Magnesium chloride supplements are commonly found in tablet and capsule forms and are used to help increase magnesium levels in the body. Magnesium is an important nutrient and is responsible for many processes in the body including regulation of blood sugar and blood pressure.

Using Section 20 in the Data Booklet, explain why the value for the enthalpy of hydration for the fluoride ion is more negative than that for the chloride ion.

The enthalpy of solution for magnesium chloride was measured in a lab as -73 kJ mol^-1
Using Sections 18 and 20 in the Data Booklet and showing your working, determine the enthalpy of hydration of magnesium ions, ΔH^θ_hyd Mg²⁺

Calculate the percentage error for your value for the enthalpy of hydration of magnesium ions, ΔH^θ_hyd Mg²⁺_, and the value given in section 18 in the Data Booklet.

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.

The enthalpy of hydration of anhydrous magnesium sulfate, MgSO₄ (s), is difficult to determine experimentally, but can be determined by using a Hess’s Law cycle.

A group of students decided to measure the enthalpy of hydration of anhydrous magnesium sulfate, MgSO₄ (s), by dissolving 3.05 g into 50.0 cm³ of water and recording the maximum temperature change. They calculated the value to be -85 kJ mol^-1.

The same group of students repeated the experiment using hydrated magnesium sulfate, MgSO₄.7H₂O (s), and calculated the enthalpy change to be +16 kJ mol^-1.

Using the student’s data, draw a Hess’s Law cycle to determine the enthalpy of hydration of solid anhydrous magnesium sulfate, MgSO₄ (s).

Determine a value for the enthalpy of hydration of anhydrous magnesium sulfate MgSO₄ (s).

A student measured the energy change when 1.35 g of zinc was added to 50 cm³ of 0.5 mol dm^-3 copper sulfate, CuSO₄ (aq), solution. The initial temperature of 21 ^oC was recorded before the addition of the zinc and a temperature reading was taken every 30 seconds.

Use the graph to determine the overall temperature change for the reaction

Calculate the enthalpy change for the reaction in kJ mol^-1.

Calculate the percentage error between your value for the enthalpy change of reaction and the literature value of -217 kJ mol^-1. Give your answer to two significant figures. 

Explain why your calculated value for the enthalpy change of reaction is different from the literature value of -271 kJ mol^-1. 

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 18 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 section 8 and 18 to calculate the enthalpy of atomisation for chlorine in strontium chloride.

The enthalpy of hydration becomes less exothermic as you go down Group 1. Explain why the enthalpy of hydration of Group 1 ions is negative.

Explain why the enthalpies of hydration become less negative as you go down Group 1. 

A Group 1 bromide has an enthalpy of solution, ΔH^Ө_sol , of 19.87 kJ mol^-1 and the lattice enthalpy, ΔH^Ө_latt , is 691 kJ mol^-1 . Use section 20 of the data booklet to identify the Group 1 ion, showing your working. 

The same Group 1 metal from part c) forms an ionic lattice with another halide ion. This new ionic compound has a larger value for lattice enthalpy, ΔH^θ_latt. Suggest a formula for the new ionic lattice and justify your answer. 

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 8 in the data booklet and the information in the table to calculate the lattice enthalpy of aluminium oxide.

Aluminium oxide is insoluble in water, but sodium oxide is soluble. Explain why there is no enthalpy of solution data for sodium oxide.

A student carried out a calorimetry experiment using 12.41 g of ammonium chloride and 12.50 cm³ of water. The temperature decreased from 23.7 °C to 17.3 °C. Construct a dissolution cycle for this reaction. 

The enthalpy change for the hydration, ΔH^θ_hyd , of the ammonium ion is -331 kJ mol^-1. Use sections 19 and 20 and your answer to part a) to calculate the lattice enthalpy, ΔH^θ_latt , of ammonium chloride.

Use sections 1, 2  and 6 in the data booklet to determine the energy change, ΔH_r , in kJ mol^-1 , for the calorimetry experiment outlined in part a).

Determine the percentage uncertainty in the student’s temperature change using a thermometer with an uncertainty of ±0.1° C.