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Chemistry - Module 5 > 22 - enthalpy and entropy > Flashcards

22 - enthalpy and entropy Flashcards

1
Q

define lattice enthalpy

A

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

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2
Q

is lattice enthalpy positive or negative and why

A

negative

lattice enthalpy involves bond formation - exothermic

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3
Q

describe each step in route 1 + 2 - is it exothermic or endothermic

A

Route 1
* Forming gaseous atoms – endothermic as bond breaking
* Forming gaseous ions – endothermic
* Making solid ionic lattice - Lattice enthalpy + exothermic

Route 2
* Enthalpy change of formation – exothermic

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4
Q

define Standard enthalpy change of formation

A
  • Enthalpy change that takes place when one mole of a compound is formed from its element under standard conditions, with all reactants and products in standard states
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5
Q

is enthalpy change of formation exothermic or endothermic

A

exothermic - bond making

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6
Q

define Standard enthalpy change of atomisation

A
  • 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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7
Q

is enthalpy change of atomisation exothermic or endothermic

A

Always endothermic – bonds are broken to form gaseous atoms

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8
Q

what do you need to be careful of when using enthalpy change of atomisation

A
  • When in standard state the bond enthalpy refers to bond being broken
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9
Q

define first ionisation energy

A
  • 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

is first ionisation energy exothermic or endothermic

A

Endothermic – energy required to overcome attraction between electron and nucleus

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11
Q

define first electron affinity

A
  • 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

is first electron affinity endothermic or exothermic

A
  • Exothermic – electron being added is attracted towards nucleus
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13
Q

is second electron affinity exothermic or endothermic

A

ENDOTHERMIC = electron being gained by negative ion which repels

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14
Q

construct this cycle

A
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15
Q

what do you need to be careful of when constructing a cycle with 2+ charges

A
  • Need first and second ionisation energies of 2+ ion
  • Two atoms of 1+ charge atom throughout – needs to be doubled
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16
Q

construct this cycle and calculate lattice enthalpy of magnesium chloride

A
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17
Q

define enthalpy change of solution

A
  • Enthalpy change that takes place when one mole of a solute dissolves in a solvent
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18
Q

enthalpy change of solution - what happens if solvent is water

A

ions from lattice end up surrounded by water molecules as aq ions because partial charges of water attracted towards ions

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19
Q

is enthalpy change of solution endothermic or exothermic

A

can be both

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20
Q

describe the experimental method of finding the enthalpy change of solution for KCl

A

Mass in equation – mass of substance changing temp

Number of moles of solute – definition says 1 mole so need to divide energy change by the number of moles

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21
Q

how does dissolving work

A
  • Ionic lattice breaks up forming separate gaseous ions

Opposite energy change to lattice energy

  • Water molecules are attracted to and surround the ions

Separate gaseous ions interact with polar water moles to form hydrated aq ions = enthalpy change of hydration

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22
Q

define enthalpy change of hydration

A
  • Enthalpy change that accompanies the dissolving of gaseous ions in water to form 1 mole of aqueous ions
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23
Q

draw a born harber cycle connecting lattice enthalpy, enthalpy change of hydration and enthalpy change of solution

A
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24
Q

calculate lattice enthalpy of sodium chloride

A
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25
Q

general properties of ionic compounds

A
  • High melting + boiling point
  • Soluble in polar solvents
  • Conduct electricity when molten or in aq solution
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26
Q

Factors affecting lattice enthalpy

A

ionic size

ionic charge

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27
Q

how does ionic size affect lattice enthalpy

A
  • Increase radius
  • Charge is more spread out over ion
  • Ions are further apart = bigger distance between centre of ions
  • Electrostatic forces of attraction between ions in lattice are weaker
  • Weaker forces = less exothermic = less negative lattice enthalpy

less energy released upon its formation from gaseous ions

28
Q

how does ionic charge affect lattice enthalpy

A
  • Increase ionic charge
  • Greater charge density
  • Stronger electrostatic forces of attraction
  • More negative = more exothermic

more energy released upon its formation from gaseous ions

29
Q

factors affecting enthalpy change of hydration

A

ionic size

ionic charge

30
Q

how does ionic size affect enthalpy change of hydration

A
  • Radius increases
  • Charge density decreases
  • Weaker attraction
  • Less energy released when they become hydrated = less exothermic
31
Q

how does ionic charge affect enthalpy change of hydration

A
  • Increase charge
  • Greater charge density
  • Stronger attractions between water molecules and ions in solution
  • More energy released when they become hydrated – more exothermic
32
Q

correlation between lattice enthalpy and melting point

A
  • More exothermic // more negative lattice enthalpy = higher melting point
33
Q

what must be overcome to dissolve ionic compound in water

A

attraction between ions in ionic lattice must be overcome

34
Q

how much energy is needed to overcome attraction between ions in an ionic lattice

A
  • Requires an amount of energy equal to the lattice enthalpy
35
Q

how to use enthalpy changes to predict solubility

A
  • If sum of hydration enthalpies greater than magnitude of lattice enthalpy – overall enthalpy change will be exothermic and compound SHOULD dissolve
36
Q

what is entropy

A
  • Measure of how disordered a system is
37
Q

units of entropy

38
Q

the greater the entropy value…

A

greater that energy is spread out per Kelvin per mole

39
Q

what state has the smallest and largest entropy values in general

A

Solids – smallest entropies

gases - greatest

40
Q

symbol for entropy

41
Q

what does 0K mean for entropy

A

no energy and all substances have entropy value of zero

42
Q

what happens if a system becomes more or less random

A
  • If system becomes more random – energy spread out more – positive entropy change
  • If system becomes less random – energy more concentrated – negative entropy change
43
Q

how does going from solid - liquid - gas affect entropy change

A
  • Entropy increases as goes from solid – liquid – gas
  • Melting / boiling increase randomness of particles
  • Energy spread out more + positive entropy change
44
Q

what type of reaction can cause an increase in entropy

A

those that increase the number of gas molecules

45
Q

if number of gas molecules increase or decrease…

A
  • If number of gas molecules decrease – negative entropy change
  • If number of gas molecules increase – positive entropy change
46
Q

predict entropy change of this reaction

47
Q

what is standard entropy

A
  • Entropy of one mole of a substance under standard conditions
48
Q

is standard entropy always positive or negative

A

always positive

49
Q

how to calculate entropy changes

A
  • must take into account number of moles - multipy entropy change for a species by number of moles in equation
50
Q

when does a reaction occur

A

if products have a lower overall energy than reactants

51
Q

what is free energy

A

overall change in energy in a chemical reaction

52
Q

symbol for gibbs free energy

53
Q

what makes up gibbs free energy

A
  • Enthalpy change – the heat transfer between chemical system and surroundings
  • Entropy change at the temp of the reaction T^S – dispersal of energy within the chemical system itself
54
Q

what is the gibbs equation

55
Q

whats the condition for feasibility

A

if the products have a lower free energy than the reactants

change in G must be negative

56
Q

determine feasibility

58
Q

endothermic feasibility

59
Q

limitations on these feasibility predictions

A
  • Does not take into account kinetics / rate of reaction – may occur at an unobservable rate
  • might be a large energy barrier (Ea) which the reacting species have to overcome before a reaction can occur
  • some reactions are feasible since ∆G is negative, but kinetically not feasible since it just occurs too slowly
  • Only predict feasibility under standard conditions
60
Q

whats the difference between standard enthalpy change of formation + lattice enthalpy

A

lattice enthalpy must start from gaseous ions to form an ionic compound

formation must start from elements in standard states to form compound

61
Q

explain whether water or ice would have the greater entropy

A

Water will have a greater entropy

particles of water are able to move more freely compared to ice

62
Q

units of entropy

A

J mol-1 K-1

63
Q

if the question has not given you enthalpy values - how can you tell if a reaction is feasible

A

at room temp - T will be positive

S value - look for number of gaseous moles on each side

H - is it exothermic (negative ) or endothermic (positive)

combustion - exo
neutralisation - endo

64
Q

how to calculate enthalpy change

A

products minus reactants

65
Q

assumptions made when calculating feasibility

Chemistry - Module 5 (7 decks)

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