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A Level Chemistry > Module 5: Chapter 22 - Enthalpy and Entropy > Flashcards

Module 5: Chapter 22 - Enthalpy and Entropy Flashcards

1
Q

What is Lattice enthalpy?

A

The enthalpy change that accompanies the formation of one mole of an ionic compounds from its gaseous ions under standard conditions

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2
Q

What is the shorthand writing of Lattice enthalpy change?

A

ΔʟᴇH

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3
Q

What type of energy change is lattice enthalpy?

A

Exothermic. Lattice enthalpy will always be exothermic as it involves the formation of ionic bonds, therefore it will always have a negative value

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4
Q

How can you determine lattice enthalpy?

A

Lattice enthalpy cannot be directly measured, therefore it must be calculated using other data in a born-haber cycle

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5
Q

What is a born-haber cycle?

A

A type of energy cycle used to analyse reaction energies

example of a born haber cycle

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6
Q

What is the standard enthalpy change of formation, ΔfH?

A

The enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states

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7
Q

What is the standard enthalpy change of atomisation, ΔₐₜH?

A

The enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions

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8
Q

What type of process is the enthalpy change of atomisation?

A

It is always an endothermic process as bonds are broken to form gaseous atoms. Therefore, it is always a positive value

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9
Q

What is first ionisation energy, ΔɪᴇH?

A

The enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

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10
Q

What type of process is ionisation energy?

A

It is always an endothermic process as energy is required to overcome the attraction between positive nuclei and negative electon. Therefore, it always has a positive value

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11
Q

What is first electron affinity, ΔᴇᴀH?

A

The enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

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12
Q

What type of process is first electron affinity?

A

First Electron affinity is always an exothermic process

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13
Q

How does second electron affinity differ from first electron affinity?

A

First electron affinity is exothermic, however second electron affinity (and all subsequent electron affinitys) are endothermic. This is because the second electron is being gained by a negative anion which repels the electron away. Therefore, energy must be put in to force the negatively-charged electron onto the negative ion

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14
Q
A

Lattice Enthalpy = -790 kJ mol⁻¹

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15
Q
A

Lattice enthalpy = -2528 kJ mol⁻¹

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16
Q

What is the standard enthalpy change of solution, ΔₛₒₗH?

A

The standard enthalpy change of solution is the enthalpy change that takes place when one mole of a solute dissolves in a solvent

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17
Q

What does “+ aq” mean in a chemical equation?

A

add excess H₂O(l)

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18
Q

How can you experimentally determine the enthalpy change of solution?

A
  1. Weigh out a sample of the ionic solid
  2. Measure out 25cm³ of water and place thermometer in the water
  3. Pour all of the solid into the water, stir the mixture and record the temperature change until it no longer increases
  4. Calculate the energy change using the equation “q = mcΔt” where m is the sum of the mass of the water and solid
  5. Calculate number of moles of ionic substance
  6. Calculate enthalpy change of solution by dividing energy by number of moles
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19
Q

What are the 2 stages of dissolving an ionic substance in water?

A
  • The ionic lattic breaks up
  • Water molecules are attracted to, and surround, the ions
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20
Q

Describe the energy changes involved with the dissolving process (enthalpy change of solution)

A
  1. The ionic lattice is broken up, forming separate gaseous ions. This is the opposite process to lattice enthalpy. This process is always endothermic
  2. The separate gaseous ions interact with polar water molecules to form hydrated aqueous ions. The energy change involved in this is known as enthalpy change of hydration. This process is always exothermic
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21
Q

What is enthalpy change of hydration, ΔhydH?

A

The enthalpy change of hydration is the enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions

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22
Q

Is enthalpy change of solution (ΔₛₒₗH) endothermic or exothermic?

A

It depends on the relative sizes of the reverse lattice enthalpy (always endothermic) and the enthalpy change of hydration (always exothermic)

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23
Q

What factors affect lattice enthalpy and hydration enthalpy?

A
  • Ionic size
  • Ionic charge
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24
Q

How does ionic size effect lattice enthalpy?

A

As ionic radius increases, the attraction between ions decreases, therefore lattice enthalpy becomes less negative (decreases in magnitude)

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25
How does ionic charge affect lattice enthalpy?
As ionic charge increases, the attraction between ions increases, therefore lattice enthalpy becomes more negative (increases in magnitude)
26
Describe the effect of ionic charge and size on attraction across period 3 from sodium to aluminium:
From sodium to aluminium: * Increasing charge gives more attraction * Decreasing size gives more attraction These are 2 **supporting** effects
27
Describe the effect of ionic charge and size on attraction across period 3 from chlorine to sulfur:
From chlorine to sulfur: * Increasing charge gives more attraction * Increasing size gives less attraction These are 2 **opposing** effects
28
How can lattice enthalpy be used to predict melting points?
**Usually** the more exothermic the lattice enthalpy is, the higher the melting point. However it can also depend on other factors such as the packing of ions in an ionic lattice
29
How does ionic size affect hydration enthalpy?
As ionic radius increases, the attraction between the ion and the water molecule decreases, so hydration energy becomes less negative
30
how does ionic charge affect hydration enthalpy?
As ionic charge increases, the attraction with water molecules increases, therefore hydration energy becomes more negative
31
How can you predict solubility?
To dissolve an ionic compound in water, the attraction between the ions in the ionic lattice must be overcome. This requires a quantity of energy equal to the lattice enthalpy. Therefore, if the sum of the hydration enthalpies is larger than the magnitude of the lattice enthalpy, the overall energy change will be exothermic and the compound *should* dissolve. However temperature and entropy can also affect this and therefore some compounds are soluble with endothermic enthalpy changes
32
What is entropy?
The natural tendency for energy to spread out rather than to be concentrated in one place. The greater the entropy, the greater the dispersal of energy and the greater the disorder
33
What are the units of entropy?
J K⁻¹ mol⁻¹
34
In general, how does entropy differ between states of matter?
* Solids have the smallest entropies * Liquids have greater entropies * Gases have the greatest entropies
35
What can entropy be used to explain?
* A gas spreading out through a room * Heat from a fire spreading through a room * Ice melting in a hot room In all of these examples energy is being dispersed and becoming more spread out
36
Explain entropy values at and above 0K
* At 0K a substance would have an entropy value of 0 * Above 0K energy becomes dispersed amongst the particles and all substances have a positive **entropy** value
37
Explain the entropy change (ΔS) if a system changes to become more random:
If a system changes to become more random, energy can be spread out more, therefore there will be a positive entropy change
38
Explain the entropy change (ΔS) if a system changes to become less random:
If a system changes to become less random, energy can be spread out less, therefore there will be a negative entropy change
39
How can you predict entropy changes in physical changes/chemical reactions?
The side of the equation with a greater number of moles of gas has greater entropy
40
Explain how melting/boiling causes an increase in entropy:
Melting and boiling increases the randomness of particles, energy is being spread out more and therefore there is an increase in entropy
41
What is the Standard entropy (Sθ) of a substance?
The entropy of one mole of a substance, under standard conditions of 100kPa and 298K
42
How can you calculate entropy changes?
ΔSθ = ΣSθ(products) - ΣSθ(reactants)
43
Can entropy have a negative value?
No, entropy values are always positive (except 0 at 0K). However entropy **changes** can be negative
44
What is the feasibility of a reaction?
The term used to describe whether a reaction is able to happen and is energetically feasible
45
What does spontaneous mean?
Energetically feasible
46
What is free energy change ΔG?
The overall energy change during a chemical reaction
47
What does free energy change consist of?
* The enthalpy change ΔH. THis is the heat transfer between the chemical system and the surroundings * The entropy change at the temperature of the reaction TΔS. This is the dispersal of energy within the chemical system itself
48
How can you calculate the free energy change?
The Gibbs' free energy equation
49
What is the Gibb's equation?
ΔG = ΔH - TΔS
50
What are the conditions for energetic feasibility of a reaction?
* There must be a decrease in free energy * ΔG < 0
51
What are the units of ΔS **in the Gibb's equation**?
**They must be *changed* to kJ K⁻¹ mol⁻¹**. Normally the units are J K⁻¹ mol⁻¹, however they must be changed to kJ K⁻¹ mol⁻¹ to match the kJ in the ΔH
52
What are the limitations of predictions made for feasibility based on ΔG?
Despite indacting thermodynamic feasibility, **it does not account for the kinetics or rate of reaction**. Therefore there may be an extremely high activation energy and the rate of reaction may be extremely low so it may not seem like the reaction is occuring
53
Calculate the lattice enthalpy for sodium bromide
-752 kJ mol⁻¹
54
Calculate the lattice enthalpy of sodium oxide
-2520 kJ mol⁻¹
55
Calculate the lattice enthalpy of calcium fluoride
-2611 kJ mol⁻¹
56
Explain the difference between the lattice enthalpies of NaCl and MgCl₂
Mg²⁺ ions are smaller and have a greater charge than Na⁺, therefore they have a greater charge density. As a result, there is a greater attraction between ions in MgCl₂ compared to NaCl. Therefore, the lattice enthalpy is more exothermic
57
Explain the difference between the lattice enthalpies of NaBr and KCl
Na⁺ ions are smaller than K⁺ ions and therefore have a greater attraction, however Cl⁻ ions are smaller than Br⁻ ions, therefore there are competing factors and it is hard to determine which lattice enthalpy will be more exothermic
58
Explain the difference between the lattice enthalpies of MgO and MgF₂
F⁻ ions are smaller than O²⁻ ions but have a smaller charge, therefore there are competing factors and it is hard to determine which lattice enthalpy will be more exothermic
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A Level Chemistry (31 decks)

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