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A Level - Chemistry > 13: Lattice energy and Entropy > Flashcards

13: Lattice energy and Entropy Flashcards

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1
Q

Define lattice energy

A

The energy change when one mole of an ionic solid is formed from its gaseous ions

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2
Q

Define enthalpy of formation

A

Enthalpy change when one mole of a substance is formed from its constituent elements in their standard states in standard conditions
Exothermic (normally)

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3
Q

Define enthalpy of combustion

A

Enthalpy change when one mole of a substance undergoes complete combustion in oxygen with all substances in the standard states in standard conditions
Exothermic

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4
Q

Define enthalpy of neutralisation

A

Enthalpy change when 1 mole of water is formed from in a reaction between an acid and an alkali in standard conditions
Exothermic

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5
Q

Define first ionisation enthalpy

A

Enthalpy change when each atom in one mole of gaseous atoms loses one electron to form one mole of gaseous 1+ ion
Standard conditions
Endothermic

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6
Q

Define second ionisation enthalpy

A

Enthalpy change when each ion in one mole of gaseous 1+ ions loses one electron to form one mole of gaseous 2+ ions
Standard conditions
Endothermic

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7
Q

Define first electron affinity

A

Enthalpy change when each atom in one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions
Standard conditions
Exothermic (non-metals)

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8
Q

Define second electron affinity

A

Enthalpy change when each ion in one mole of gaseous 1- ions gain one electron to form one mole of gaseous 2- ions
Standard conditions
Endothermic (adding -ve electron to a -ve ion)

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9
Q

Define enthalpy of atomisation

A

Enthalpy change when one mole of gaseous atoms is produce from an element in its standard state
Standard conditions
Endothermic

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10
Q

Define enthalpy of solution

A

Enthalpy change when one mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are well separated and don’t interact with one another
Standard conditions
Varies between exo and endo

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11
Q

Define hydration enthalpy

A

Enthalpy change when one mole of gaseous ions become hydrated (dissolved in water)
Standard conditions
Exothermic

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12
Q

Define bond dissociation enthalpy

A

Enthalpy change when one mole of covalent bonds is broken in the gaseous state
Standard Conditions
Endothermic

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13
Q

Define lattice enthalpy of formation

A

Enthalpy change when one mole of a solid ionic compound is formed from its constituent ions in the gas phase
Standard conditions
Exothermic

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14
Q

Define lattice enthalpy of dissociation

A

Enthalpy change when one mole of a solid ionic compound is broken up into its constituent ions in the gas phase
Standard conditions
Endothermic

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15
Q

What is the relationship between lattice enthalpy of formation and dissociation?

A

Lattice enthalpy of formation = - (lattice enthalpy of dissociation)

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16
Q

Define enthalpy of vaporisation

A

Enthalpy change when one mole of a liquid is turned into a gas
Standard conditions
Endothermic

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17
Q

Define enthalpy of fusion

A

Enthalpy change when one mole of a solid is turned into a liquid
Standard conditions
Endothermic

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18
Q

What is the cycle for enthalpy of solution calculations?

A

Enthalpy of solution = Ionic solid -> Dissolved ions
Lattice enthalpy of formation = Gas ions -> ionic solid
Hydration enthalpies = Gas ions -> Dissolved ions

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19
Q

What is the equation for enthalpy of solution calculations?

A

Enthalpy of solution = -(lattice enthalpy of formation) + hydration enthalpies (both ions)

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20
Q

What does the lattice enthalpy of a compound indicate?

A

Strength of ionic bonding - larger the magnitude of lattice enthalpy, the stronger the bonding

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21
Q

Which ions will have higher lattice enthalpy due to stronger attractions?

A

Smaller ions, and ions with higher charge

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22
Q

How is the experimental lattice enthalpy of formation found?

A

Born-Haber cycle

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23
Q

What are the stages to take elements in normal states to become a ionic compound?

A

Atomisation of metal and non-metal (endo)
Ionisation of metal atoms (endo)
Electron affinity of non-metal atoms (exo)
Lattice enthalpy of formation (exothermic)

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24
Q

What is experimental lattice enthalpy?

A

Lattice enthalpy calculate using a Born-Haber cycle
All other enthalpy changes found experimentally
This is the real value

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25
Q

What is theoretical lattice enthalpy?

A

Lattice enthalpy calculated using an equation considering: size, charge, etc.
Assumes the structure is perfectly ionic
Not the real value

26
Q

What is the real lattice enthalpy value?

A

The experimental lattice enthalpy

27
Q

Why is the theoretical lattice enthalpy not the real value?

A

Often distortion of the ions (polarisation), meaning the ions are not perfectly spherical
Therefore it is not perfectly ionic, said to have covalent character

28
Q

How do cations effect ionic structure?

A

Positive charge with a high charge density are very good at distorting negative ions

29
Q

How do anions effect ionic structure?

A

Negative ions that are large or highly charged are easier to distort

30
Q

What is the characteristics of ionic compounds which have covalent character?

A

Low solubility in water and lower melting points and electrical conductivity

31
Q

What is the general trend of the difference between experiment and theoretical lattice enthalpy values?

A

Larger the difference between the values, the greater the covalent character

32
Q

What is polarisation?

A

The ability of a cation to draw electrons from an anion closer to the cation
Anion is no longer spherical

33
Q

What does the enthalpy of a solution have to be for a solid to dissolve?

A

Exothermic (negative)

However if slightly positive may react due to entropy

34
Q

What is the total hydration enthalpy change when dissolving ionic solids?

A

Enthalpy of hydration of the cation + anion

35
Q

What are the two steps of dissolving ionic solids?

A

Ions in solid become separated from each other, and turn into gaseous ions
Gaseous ions are hydrated, as ions form bonds with the solvent water molecules

36
Q

What bonding forms between water and ions?

A

Cations (+) attracted to 𝛿- oxygen atoms

Anions (-) attracted to 𝛿+ hydrogen atoms

37
Q

How does the solubility of group 2 sulphates and carbonates change going down the group and why?

A

Solubility decreases down the group
Radius of cation increases , hydration enthalpy becomes significantly less exothermic
Lattice enthalpy decreases, this is NOT significant
Overall enthalpy of solution becomes less exothermic

38
Q

How does the solubility of group 2 hydroxides change going down the group and why?

A

Solubility increases down the group
Radius of cation increases, hydration enthalpy NOT significantly less exothermic
Lattice enthalpy decreases, significantly
Overall the enthalpy of solution becomes more exothermic

39
Q

What means some endothermic reactions occur at room temperatures?

A

Entropy must be increasing

40
Q

Define entropy

A

A measure of the disorder of a system

41
Q

What is the natural direction of change of entropy?

A

Increasing total entropy (+ve entropy change)

42
Q

What entropy will a solid very ordered system have?

A

Very low entropy

43
Q

What are the entropies of different states?

A

Solids - low
Liquids - higher
Gas - much higher

44
Q

What explains the entropies of different states?

A

Gases have many different combinations which they can be found as they are not tightly packed

45
Q

What happens to entropy as a molecule becomes more complex and why?

A

Increases in entropy

Due to more combinations which can occur due to more molecules and more electrons so more disordered

46
Q

Why do the size of the atoms affect entropy?

A

Larger atoms have more entropy due to more different configurations which can occur

47
Q

What is ∆S system?

A

Entropy change which occurs in the reaction

equation

48
Q

How do you calculate ∆S system?

A

Entropy of products - reactants

49
Q

What is ∆S of surroundings?

A

The entropy change of the surroundings when a reaction occurs (doesn’t include the changes of the molecules in the reaction)

50
Q

How do you calculate ∆S surroundings?

A

-∆H / T
∆H - enthalpy change of the reaction
T - Temperature in kelvin

51
Q

What is the equation for ∆S total?

A

∆S system + ∆S surroundings

52
Q

What does ∆S total have to be for a reaction to be feasible?

A

∆S total must be positive

53
Q

Why do liquids have slightly higher entropy values than solids?

A

Atoms/molecules are free to move meaning more combionations

54
Q

What are the standard units for ∆S?

A

J/K/mol

55
Q

What does spontaneous mean in a thermodynamics context?

A

A reaction in spontaneous if it leads to an overall increase in entropy
(which makes it feasible)

56
Q

What is Gibbs free energy equation?

A

ΔG = ΔH - TΔS system

57
Q

What should Gibbs free energy be for a reaction to be feasible?

A

The Gibbs free energy value has to be negative for a reaction to be feasible

58
Q

What is the equation for calculating temperature in gibbs?

A

T = ΔH /ΔS system

59
Q

What is the equation relating ΔS total and the equilibrium constant?

A

ΔS total = R lnK

60
Q

Predict temp change on equilibrium constant if ΔH = 0, assume ΔS system doesn’t change with temp

A 
ΔH = 0
Therefore ΔS surrounding = 0 as -ΔH / T
ΔS total = ΔS system
ΔS total = R lnK
ΔS system and total not affected by temp, so K doesn't change

A Level - Chemistry (26 decks)

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