Year 12 Chemistry, Study Guides, Projects, Research of Chemistry

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Year 12 Chemistry

Ark Globe Academy

Remote Learning Pack

Phase IV

Monday 8 June- Friday 19 June

Year 12 Chemistry

Session Title Work to be completed Resource provided Outcome On-Line Support 1 Recap of AS Energetics Read through content to recap AS Energetics and apply to related recap questions in task 1 Work pack IV Page(s) 3 – 8 Complete: Task 1 Session 1 Recap of AS Energetics (VLE) 2 Definitions Review the table of key definitions and use these to complete task 2. Work pack IV Page(s) 9 – 10 Textbook Page(s): 258 – 260 Complete: Task 2 3 MS Teams Lesson Wednesday Live lesson: Enthalpy of Solution Calculations Work pack IV Pages(s) 11 – 13 Textbook page(s): 267 – 268 Complete: 1 - DO NOW 2 - Checking for Understanding Question(s) 3 - Application Question(s) 4 Enthalpy of Solution Calculations Complete task 3 questions using information from previous session Work pack IV Pages(s) 14 Textbook page(s): 267 – 268 Complete: Task 3 Session 4 Enthalpy of Solution Calculations (VLE) 5 MS Teams Lesson Friday Live lesson: Born Haber Cycle Work pack IV Pages(s) 15 – 17 Textbook page(s): 261 – 266 Complete: 1 - DO NOW 2 - Checking for Understanding Question(s) 3 - Application Question(s) 6 Born Haber Cycle Calculations Complete task 4 questions using information from previous session Work pack IV Pages(s) 18 Textbook page(s): 261 – 266 Complete: Task 4 Session 6 Born- Haber Cycle Calculations (VLE) 7 Application of New Theory Complete a series of PPQs on the theory you have covered since session 2 Work pack IV Pages(s) 19 – 22 Complete: Past paper exam questions, Q. 1 to Q. 5 8 MS TEAMS Lesson Wednesday Live lesson: Comparing Experimental and Theoretical Enthalpy Values Work pack IV Pages(s) 23 - 24 Textbook page(s): 267 – 268 Complete: 1 - DO NOW 2 - Checking for Understanding Question(s) 3 - Application Question(s)

Session 1 - Recap of AS Energetics

1) What is enthalpy?

• What is enthalpy? It is a measure of the heat content of a substance
• Enthalpy change (ΔH) = Change in heat content at constant pressure
• Standard conditions (ΔHº) = 100 kPa and a stated temperature [Note that the symbol for standard conditions should be a circle with a horizontal line through it, but it is safer here to put a circle with no line as otherwise the character that may appear on the computer the document is from may substitute some random symbol!]

2) Reaction profiles

3) Standard enthalpy change of reaction ( Δ rH ° ) (“enthalpy of reaction”)

This is the enthalpy change for a reaction with the quantities shown in the chemical equation. This means that the value should always be quoted along with the equation. In this example, the second equation contains half the molar quantities of the first and so the (^) rH value is half as much. e.g. H 2 SO 4 (aq) + 2 NaOH(aq) → Na 2 SO 4 (aq) + 2 H 2 O(l) ΔrH° = – 114.2 kJ mol-^1 ½ H 2 SO 4 (aq) + NaOH(aq) → ½ Na 2 SO 4 (aq) + H 2 O(l) ΔrH° = – 57.1 kJ mol-^1 the value of – 114.2 kJ mol-^1 in the first equation means that 114.2 kJ of heat energy is released when 1 mole of H 2 SO 4 reacts with 2 moles of NaOH. the value of – 57.1 kJ mol-^1 in the second equation means that 57.1 kJ of heat energy is released when ½ mole of H 2 SO 4 reacts with 1 mole of NaOH.

4) Standard enthalpy change of formation ( Δ fH ° ) (“enthalpy of formation”)

Enthalpy change when 1 mole of a substance is formed from its constituent elements with all reactants and products in standard states under standard conditions. e.g. CH 4 (g) C(s) + 2 H 2 (g) → CH 4 (g) H 2 O(l) H 2 (g) + ½ O 2 (g) → H 2 O(l) NH 3 (g) ½ N2(g) + 3/2 H2(g) → NH3(g) C 2 H 5 OH(l) 2 C(s) + 3 H2(g) + ½ O2(g) → C2H5OH(l) CH 3 Br(l) C(s) + 3/2 H2(g) + ½ Br2(l) → CH3Br(l) Na 2 O(s) 2 Na(s) + ½ O2(g) → Na2O(s) Note: re ΔfHº of an element in its standard state = 0 by definition

5) Standard enthalpy change of combustion (∆cH°) (“enthalpy of combustion”)

Enthalpy change when 1 mole of a substance is completely burned in oxygen with all reactants and products in standard states under standard conditions. e.g. CH4(g) CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) H2(g) H2(g) + ½ O2(g) → H2O(l) C2H6(g) C2H6(g) + 3½ O2(g) → 2 CO2(g) + 3 H2O(l) C2H5OH(l) C2H5OH(g) + 3 O2(g) → 2 CO2(g) + 3 H2O(l) Na(s) Na(s) + ¼ O2(g) → ½ Na2O(s) C6H14(l) C6H14(g) + 9½ O2(g) → 6 CO2(g) + 7 H2O(l)

2) Calculations involving enthalpies of combustion (“Type 2 questions”)

• Questions that involve enthalpies of combustion can usually be done using the cycle shown.
• The reaction involved across the top is often an enthalpy of formation (from elements to a compound).
• The sum of the clockwise arrows equals the sum of the anticlockwise arrows.
• Be careful when drawing your cycle to ensure that arrows are going in the right direction and the number of moles is correct.
• If you use a cycle like this, there is no need to worry about getting the number of oxygen molecules in the downward arrows.

3) Calculations involving bond enthalpies (“Type 3 questions”)

• Bond enthalpy is the enthalpy change to break one mole of covalent bonds in the gas phase.
• For most bonds (e.g. C-H, C-C, C=O, O-H, etc.) the value for the bond enthalpy is an average taken from a range of molecules as the exact value varies from compound to compound. For some bond enthalpies (e.g. H-H, H-Cl, O=O, etc) the value is exact as only one molecule contains that bond.
• Mean bond enthalpy is the enthalpy change to break one mole of covalent bonds in the gas phase averaged over several compounds.
• Enthalpies of reaction that have been calculated using mean bond enthalpies are not as accurate as they might be because the values used are averages and not the specific ones for that compound.
• The following cycle works for any question that involves bond enthalpies, whether to find a bond enthalpy or ΔH for a reaction.
• Remember that substances must be in the gas state before bonds are broken, and so ΔH to go to the gas state is needed for solids and liquids. (Note - ΔH vaporisation is the enthalpy change to convert a liquid to a gas)
• As with other cycles, the sum of the clockwise arrows equals the sum of the anticlockwise arrows. Be careful to ensure that arrow directions and number of moles are correct.

4) Calorimetry calculations

• The enthalpy change for a reaction can be found by measuring the temperature change in a reaction.
• The heat energy given out (or taken in) is used to heat (or cool) a known mass of water. We know that it takes 4.18 J of energy to raise the temperature of 1 g of water by 1°C (i.e. 1 K).
• The amount of energy needed to make 1 g of a substance 1°C (1 K) hotter is called the specific heat capacity (measured in J g-^1 K-^1 ).
• The following equation is then used to find the amount of heat energy give out (or absorbed).
• To find the enthalpy change in terms of J (or kJ) per mole, the following expression is needed: ( THINK kJ per mole!)
• Heat loss is a major problem with calorimetry and can lead to errors in the results. The techniques used in calorimetry are designed to reduce heat loss (one way to reduce errors from heat loss is to measure the heat capacity of the calorimeter as a whole).

Session 2 - Definitions

TASK 2 – Enthalpy Change Definitions: Write equations for reactions described below using

definitions given in the table on the previous page. 1. ΔfH of C 6 H 6 (l) ………………………………………………………………...………………………….. 2. ΔfH of CH 3 COOH(l) ……………………………………………………………………………………... 3. ΔcH of H 2 (g) ……………………………………………………………………………………………… 4. ΔcH of CH 3 COOH(l) …………………………………………………………………………………….. 5. 1st ionisation energy of aluminium ……………………………………………………………….…… 6. 2nd ionisation energy of aluminium ………………………………………………………………...… 7. 3rd ionisation energy of aluminium …………………………………………………………………… 8. 1st electron affinity of chlorine ………………………………………………………………………… 9. lattice enthalpy of formation of sodium oxide ……………………………………………………...... 10. lattice enthalpy of dissociation of aluminium oxide ………………………………………………… 11. ΔhydH of sodium ions ………………………………………………………………………………….. 12. Enthalpy of vaporisation of bromine ………………………………………………………….......... 13. ΔsolH of sodium hydroxide …………………………………………..………...........................…… 14. Enthalpy of fusion of sodium chloride ………………………………………………...............…… 15. Bond dissociation enthalpy of water ……………………………………………….....................… 16. Bond dissociation enthalpy of hydrogen ………………………………………….......…………… 17. ΔatmH of bromine ………………………………………………………...................................…… 18. Bond dissociation enthalpy of bromine ……………………………………………………........… 19. 1st electron affinity of bromine …………………………………………………….................…… 20. 2nd electron affinity of sulphur ………………………………………………………….............…

With positive ions, there may only be loose ion-dipole attractions between the δ- oxygen atoms in the water molecules and the positive ions, or there may be formal dative covalent (co-ordinate covalent) bonds. With negative ions, ion-dipole attractions are formed between the negative ions and the δ+ hydrogens in water molecules. The size of the hydration enthalpy is governed by the amount of attraction between the ions and the water molecules.

• The attractions are stronger the smaller the ion. For example, hydration enthalpies fall as you go down a group in the Periodic Table. The small lithium ion has by far the highest hydration enthalpy in Group1, and the small fluoride ion has by far the highest hydration enthalpy in Group 7. In both groups, hydration enthalpy falls as the ions get bigger.
• The attractions are stronger the more highly charged the ion. For example, the hydration enthalpies of Group 2 ions (like Mg2+) are much higher than those of Group 1 ions (like Na+). Estimating enthalpies of solution from lattice enthalpies and hydration enthalpies The hydration enthalpies for calcium and chloride ions are given by the equations: The following cycle is for calcium chloride and includes a lattice dissociation enthalpy of +2258 kJ mol-1. We have to use double the hydration enthalpy of the chloride ion because we are hydrating 2 moles of chloride ions. Make sure you understand exactly how the cycle works.

So... ΔHsol = +2258 - 1650 + 2(-364) ΔHsol = - 120 kJ mol- 1 Whether an enthalpy of solution turns out to be negative or positive depends on the relative sizes of the lattice enthalpy and the hydration enthalpies. In this particular case, the negative hydration enthalpies more than made up for the positive lattice dissociation enthalpy.

Session 5 – MS Teams Lesson – Friday:

Born Haber Cycle

The notes in this section will help supplement what you will cover in the MS Teams lesson and allow you to successfully complete the tasks in the sessions to follow.

Lattice enthalpy

• Lattice enthalpy represents the enthalpy change when the ions in one mole of a solid ionic compound are broken apart (lattice enthalpy of dissociation) or brought together (lattice enthalpy of formation).
• The lattice enthalpy of a compound is an indication of the strength of the ionic bonding – the greater the magnitude of the lattice enthalpy, the stronger the bonding.
• Generally speaking, compounds with smaller ions and/or ions with higher charges have stronger attractions and so greater lattice enthalpy. e.g. NaCl has a higher lattice enthalpy (and therefore stronger ionic bonding) than KCl as they Na+^ ion is smaller than the K+^ ion. e.g. MgCl 2 has a higher lattice enthalpy (and therefore stronger ionic bonding) than NaCl as the Mg2+^ ion has a higher charge and is smaller than the Na+^ ion. Measuring lattice enthalpy
• The lattice enthalpy of a compound can be found using a Born-Haber cycle – this value is often called the ‘experimental value’ as the data used in the Born-Haber cycle is determined by experiments.
• A Born-Haber cycle is a cycle that includes all the enthalpy changes in the formation of an ionic compound.
• In these cycles, lattice enthalpy is usually shown as lattice enthalpy of formation (with – ve value)
• this is so that an equation can be written where the enthalpy of formation of the ionic compound (not to be confused with the lattice enthalpy of formation) equals the sum of all the other enthalpy changes. Note that when drawing Born-Haber cycles:
• draw a separate step for every enthalpy change (e.g. for atomisation of the metal atoms separately from the non-metal atoms, for each individual ionisation enthalpy, for each individual electron affinity)
• second and third electron affinities are endothermic and shown be drawn going up not down
• It is best to write the numerical values of the enthalpy changes on each step (you may also be asked to write the names of each step)

Example 1 – Find the lattice enthalpy of formation of calcium oxide using a Born-Haber cycle and these enthalpy changes:

Session 6 – Born Haber Cycle Calculations

Task 4 – Born Haber Cycle Calculations

Use this data in the questions that follow. kJmol-^1 Na K Ca Al Co Cu Br I O S Cl enthalpy of atomisation +107 +90 +193 +314 +427 +112 +107 +248 +279 + 1st ionisation enthalpy +496 +418 +590 +577 +757 + 2nd ionisation enthalpy +4562 +3070 +1150 +1820 +1640 + 3rd ionisation enthalpy +6910 +4600 +4940 +2740 +3230 + 1st electron affinity (^) – 342 – 142 – 200 – 364 2nd electron affinity +844 + 1 Calculate the enthalpy of formation of potassium chloride given that the lattice enthalpy of formation of potassium chloride is – 710 kJmol-^1. 2 Calculate the lattice enthalpy of formation of sodium sulfide given that the enthalpy of formation of sodium sulfide is – 370 kJmol-^1. 3 Calculate the enthalpy of formation of calcium bromide given that the lattice enthalpy of formation of calcium bromide is – 2125 kJmol-^1. 4 Calculate the lattice enthalpy of formation of aluminium oxide given that the enthalpy of formation of aluminium oxide is – 1669 kJmol-^1. 5 Calculate the first electron affinity of iodine given that the lattice enthalpy of dissociation of calcium iodide is +2054 kJmol-^1 and its enthalpy of formation is – 535 kJmol-^1. 6 Calculate the enthalpy of atomisation of copper given that the enthalpy of formation of CuO is – 155 kJmol-^1 and its lattice enthalpy of formation is – 4149 kJmol-^1. 7 The lattice enthalpy of formation of the three possible chlorides of cobalt are given: CoCl – 700; CoCl 2 – 2624; CoCl 3 – 5350 kJmol-^1. a) Using Born-Haber cycles, calculate the enthalpy of formation of each chloride. b) Which of these chlorides is energetically stable with respect to their elements under standard conditions? c) Which compound is likely to be formed when cobalt and chlorine react under normal conditions?

Session 7 – Application of New Theory

Q1. Which equation represents the process that occurs when the standard enthalpy of atomisation of iodine is measured? A I 2 (s) → I(g) B I 2 (s) → 2I(g) C I 2 (g) → I(g) D I 2 (g) → 2I(g) (Total 1 mark) Q2. This question is about magnesium chloride. (a) Write the equation, including state symbols, for the process corresponding to the enthalpy of solution of magnesium chloride.

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(1) (b) Use these data to calculate the standard enthalpy of solution of magnesium chloride. Enthalpy of lattice dissociation of MgCl 2 = +2493 kJ mol–^1 Enthalpy of hydration of magnesium ions = – 1920 kJ mol–^1 Enthalpy of hydration of chloride ions = – 364 kJ mol–^1

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(2) (c) Solubility is the measure of how much of a substance can be dissolved in water to make a saturated solution. A salt solution is saturated when an undissolved solid is in equilibrium with its aqueous ions. Use your answer to part (b) to deduce how the solubility of MgCl 2 changes as the temperature is increased. Explain your answer.

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(3) (Total 6 marks) Q3. Thermodynamics can be used to investigate the changes that occur when substances such as calcium fluoride dissolve in water. (a) Give the meaning of each of the following terms. (i) enthalpy of lattice formation for calcium fluoride

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(2)