A-level Chemistry: 3.1.8 Thermodynamics
This content covers a variety of thermodynamic concepts related to enthalpy changes, including formation, atomisation, dissociation, ionization, electron affinity, hydration, and lattice enthalpy. It also explores the relationship between experimental and theoretical lattice enthalpies and provides insights into the dissolution of ionic compounds in water, emphasizing the role of bond formation and energy changes in these processes.
Define enthalpy change of formation (ΔfH)
Enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions
Key Terms
Define enthalpy change of formation (ΔfH)
Enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions
Write an equation representing the enthalpy change of formation of ethanol
Define bond dissociation enthalpy (ΔdissH)
Enthalpy change when one mole of a (covalent) bond is broken into (two) gaseous atoms (or free radicals)
Write an equation representing the bond dissociation enthalpy of chlorine
Define enthalpy change of atomisation of an element (ΔatH)
Enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state
Write an equation representing the enthalpy change of
atomisation of chlorine
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| Term | Definition |
|---|---|
Define enthalpy change of formation (ΔfH) | Enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions |
Write an equation representing the enthalpy change of formation of ethanol | |
Define bond dissociation enthalpy (ΔdissH) | Enthalpy change when one mole of a (covalent) bond is broken into (two) gaseous atoms (or free radicals) |
Write an equation representing the bond dissociation enthalpy of chlorine | |
Define enthalpy change of atomisation of an element (ΔatH) | Enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state |
Write an equation representing the enthalpy change of | |
Define enthalpy change of atomisation of a compound (ΔatH) | Enthalpy change when 1 mole of a compound in its standard state in converted to gaseous atoms |
Write an equation representing the enthalpy change of atomisation of sodium chloride | |
Define first ionisation energy (Δie1H) | Enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms |
Write an equation representing the first ionisation energy of magnesium | |
Define second ionisation energy (Δie2H) | Enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions |
Write an equation representing the second ionisation energy of magnesium | |
Define first electron affinity (Δea1H) | Enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms |
Write an equation representing the first electron affinity of oxygen | |
Define second electron affinity (Δea2H) | Enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions |
Write an equation representing the second electron affinity of oxygen | |
| Enthalpy change when 1 mole of aqueous ions is formed from 1 mole of gaseous ions |
Write an equation representing the enthalpy change of hydration of sodium | |
Define enthalpy change of solution (ΔsolutionH) | Enthalpy change when 1 mole of solute is dissolved in enough solvent that no further enthalpy change occurs on further dilution |
Write an equation representing the enthalpy change of solution of sodium chloride | |
Define lattice enthalpy of formation (ΔlatticeH) | Enthalpy change when 1 mole of solid ionic compound is formed from its gaseous ions under standard conditions |
Write an equation representing the lattice enthalpy of formation of magnesium chloride | Exothermic |
Define lattice enthalpy of dissociation (ΔlatticeH) | Enthalpy change when 1 mole of solid ionic compound is completely dissociated into its gaseous ions under standard conditions |
Write an equation representing the lattice enthalpy of dissociation of magnesium chloride | Endothermic |
What can you use to calculate lattice enthalpy since you can't calculate it directly? | Born-Haber cycle |
Calculate the lattice enthalpy of formation of sodium chloride given that the enthalpy of formation of sodium chloride is -411 kJ mol-1 | The lattice enthalpy of sodium chloride is –786 kJ mol⁻¹. |
Calculate the atomisation enthalpy of magnesium given that the enthalpy of formation of magnesium oxide is -548 kJ mol-1 and lattice enthalpy of formation is -3791 kJ mol-1 | The atomisation enthalpy of magnesium is +150 kJ mol⁻¹. |
Calculate the lattice enthalpy of formation of aluminium oxide given that the enthalpy of formation of aluminium oxide is -1669 kJ mol-1 | The lattice enthalpy of formation of aluminium oxide is –15,769 kJ mol⁻¹. |
What does the purely ionic model of a lattice assume? |
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Why are experimental lattice enthalpy values different to theoretical lattice enthalpy values? |
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Explain why the differences between experimental and lattice enthalpies are much bigger for magnesium halides than sodium halides |
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When a solid ionic lattice dissolves in water, what 2 things occur? |
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Explain why water molecules can bond to ions |
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Substances generally only dissolve if the energy… | released is roughly the same, or greater than the energy taken in |
Souble substances tend to have enthalpies of solution | exothermic |
Calculate the enthalpy of solution for sodium chloride given that the lattice dissociation enthalpy is 787 kJ mol-1 and the enthalpies of hydration of sodium and chloride ions are -406 and -364 kJ mol-1 respectively. | |
Calculate the enthalpy of solution for silver chloride given that the lattice dissociation enthalpy is 905 kJ mol-1 and the enthalpies of hydration of silver and chloride ions are -464 and -364 kJ mol-1 respectively. | |
Calculate the enthalpy of solution for magnesium chloride given that the lattice formation enthalpy is -2493 kJ mol-1 and the enthalpies of hydration of magnesium and chloride ions are -1920 and -364 kJ mol-1 respectively. | |
What is entropy (S)? | Measure of no. of ways that particles can be arranged & no. of ways that energy can be shared out between particles |
More disordered the particles are, the ____ (more +ve the value) entropy is | higher |
Name and describe 2 factors that affect entropy |
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Why do substances always tend towards disorder? | They're more energetically stable when there's more disorder ∴ particles will move to increase their entropy |
Explain why the reaction of sodium hydrogencarbonate with hydrochloric acid is feasible even though it's an endothermic reaction |
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During a reaction, there's an _ ____ (ΔS) between reactants and products | Entropy change |
State the equation you would use to calculate ΔS | |
State the units of entropy | J K-1 mol-1 |
Define standard entropy of a substance (S⦵) | Entropy of 1 mole of that substance under standard conditions |
What is free energy change (ΔG)? | Measure used to predict whether a reaction is feasible |
Free Energy Change When is a reaction considered feasible? | If ΔG = -ve or 0 |
Even if ΔG shows that a reaction is theoretically feasible, it might have a … or be … that you wouldn't notice it happening at all | Even if ΔG shows that a reaction is theoretically feasible, it might have a very high activation energy or be so slow that you wouldn't notice it happening at all |
Feasibility depends on _____ | Temperature |
When is ΔG always negative (i.e. reactions feasible at any temperature)? Name the 2 conditions | If reaction is exothermic (ΔH = -ve) & has +ve entropy change |
When is ΔG always postive (i.e. reactions are NOT feasible at any temperature)? Name the 2 conditions | If reaction is endothermic (ΔH = +ve) & has -ve entropy change |
What happens to the reaction's feasibility if ΔH = positive & ΔS = positive? | Reaction won't be feasible at some temps but will be at high enough temperature |
What happens to the reaction's feasibility if ΔH = negative & ΔS = negative? | Reaction will be feasible at lower temperatures but won't be feasible at higher temperatures |
Describe how you can calculate the temperature at which a reaction becomes feasible | Can find temperature when ΔG = 0 |
Suggest why the electron affinity of chlorine is an exothermic change (2) |
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State whether you would expect the value of the theoretical enthalpy of lattice dissociation for silver chloride to be greater than, equal to or less than that for silver bromide. Explain your answer. (3) |
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Explain the meaning of the term perfect ionic model (1) | Ions can be regarded as point charges (or perfect spheres) |
Suggest why the experimental value is greater than the theoretical value for the enthalpy of lattice dissociation for silver chloride (2) |
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Explain the interaction between water molecules and fluoride ions when the fluoride ions become hydrated (2) |
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Suggest why electron affinity is an endothermic process for this reaction: |
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Explain why the hydration enthaply of a fluoride ion is more negative than the hydration enthaply of a chloride ion (2) |
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By describing the nature of the attractive forces involved, explain why the value of the enthaply of hydration for a chloride ions is more negative than that for the bromide ion (3) |
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Suggest why a value for the enthaply of solution of magnesium oxide is not found in any data books. (1) | Magnesium oxide reacts with water / forms Mg(OH)2 |
Suggest why using a Born-Haber cycle produces more accurate results than using mean bond enthalpies (1) | Mean bond enthalpies are from a range of compounds |