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Chemistry Year 12/13 > Module 5.2 - Energy > Flashcards

Module 5.2 - Energy Flashcards

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1
Q

What is meant by the term ‘lattice enthalpy’?

A

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

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2
Q

What is meant by the term ‘standard enthalpy change of formation’?

A

The enthalpy change that accompanies the formation of 1 mole of a compound from its elements

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3
Q

What is meant by the term ‘first ionisation energy’?

A

The energy change that accompanies the removal of 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

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4
Q

What energy changes occur during the formation of an ionic lattice?

A
  • when oppositely charge ions attracted to one another to form giant ionic lattice, theres a large lowering of energy from strong attraction
  • although energy required to form ions is huge, the lowering of energy when forming lattice more than compensates for this
  • hence why giant ionic lattices have v strong ionic bonds and v high melting and boiling points
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5
Q

What are the features of lattice enthalpy?

A
  • ions it forms from are both gaseous
  • 1 mol of substance formed
  • enthalpy change is negative (energy given off to surroundings)
  • ionic lattice formation is exothermic
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6
Q

What does a more exothermic lattice enthalpy indicate about the strength of the ionic bonds in the lattice and therefore the melting and boiling points?

A
  • more exothermic = stronger ionic bonds = stronger electrostatic interactions
  • high mp and bp as more energy to overcome interactions present
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7
Q

What type of ion would cause the most negative lattice enthalpy and why?

A
  • small ions with large charges
  • charges cause large electrostatic forces
  • smaller ions can get closer together
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8
Q

Why can’t lattice enthalpy be directly measured, and how would you work it out instead?

A
  • it’s impossible for one mole of an ionic lattice to form from its gaseous ions
  • use a Born-Haber cycle
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9
Q

What are the key features of a Born-Haber cycle?

A
  • continuous cycle formed that can start at the elements and end at the elements
  • includes 1 step showing formation of 1 mole of solid ionic lattice from its gaseous ions; the lattice enthalpy
  • remaining steps show intermediate changes that correspond to key enthalpy changed that can be measured
  • lattice enthalpy can be calculated by Hess’ law (if a reaction can take place by more than one route and the initial and final conditions are the same, the Toal enthalpy change for each route is the same
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10
Q

What enthalpy changes usually take place during in a Born-Haber cycle (especially for lattice enthalpy)?

A
  • standard enthalpy change of formation
  • standard enthalpy change of atomisation
  • first/second ionisation energy
  • first/second electron affinity
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11
Q

Is the standard enthalpy change of formation for an ionic lattice usually endothermic or exothermic and why?

A
  • exothermic

- bonds are made

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12
Q

Give an equation for the standard enthalpy change of formation of potassium chloride.

A

K(s) + 1/2Cl2(g) –> KCl(s)

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13
Q

Is standard enthalpy change of atomisation endothermic or exothermic and why?

A

Endothermic as bonds have to be broken

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14
Q

What does ‘standard enthalpy change of atomisation’ mean?

A

One mole of gaseous atoms formed from its element in its standard state

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15
Q

Give the standard enthalpy change of atomisation of chlorine and potassium.

A

K(s) –> K(g)

1/2Cl2(g) –> Cl(g)

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16
Q

What does ‘first electron affinity’ mean?

A

One mole of gaseous 1- ions formed from gaseous atoms (each one gains an electron)

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17
Q

Give the equation for the first electron affinity of chlorine.

A

Cl(g) + e- –> Cl-(g)

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18
Q

Is first/second ionisation energy usually exothermic or endothermic and why?

A

Endothermic as electron being lost has to overcome attraction from nucleus to leave atom

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19
Q

Is first electron affinity exothermic or endothermic and why?

A
  • exothermic

- electron is attached into outer shell of an atom by nucleus (bond formed)

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20
Q

Is second electron affinity exothermic or endothermic and why?

A
  • endothermic

- electron repelled by the 1- ion and this repulsion has to be overcome

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21
Q

What equation can you use to use Hess’ law on Born-Haber cycles?

A

sum of clockwise enthalpies = sum of anticlockwise enthalpies

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22
Q

What is meant by the term ‘enthalpy change of solution’?

A

The enthalpy change that is associated with dissolving 1 mole of solute in a solvent to form an infinitely dilute solution (i.e. completely dissolved in solvent)

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23
Q

What is meant by the term ‘enthalpy change of hydration’?

A

The enthalpy change that associates dissolving 1 mole of gaseous ions in water

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24
Q

What 2 processes take place when an ionic solid dissolves?

A
  • ionic lattice breaks down

- free ions become part of the solution (hydration)

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25
Q

Is enthalpy change of solution endothermic or exothermic?

A

Can be either

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26
Q

How is the breakdown of an ionic lattice related to lattice enthalpy?

A
  • essentially opposites, breakdown of ionic lattice equals the negative of lattice enthalpy
  • process are identical but reverse
  • enthalpy change is same but different sign
  • lattice enthalpy is exothermic (-ve) and breakdown is endothermic (+ve)
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27
Q

What factors affect lattice enthalpy?

A
  • size of ions involved
  • charges on the ions
  • ionic bond strength (dependent on ionic size and charge)
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28
Q

How does ionic size affect lattice enthalpy?

A
  • smaller ions (smaller atomic radius) can get closer together
  • attract more strongly
  • more exothermic lattice enthalpy (more -ve)
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29
Q

How does the charge of an ion affect lattice enthalpy?

A
  • higher charges causes greater electrostatic attraction

- more exothermic (more -ve)

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30
Q

What type of ions would cause the most exothermic lattice enthalpy?

A

very small, highly charged

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31
Q

What type of solvent do ionic solids dissolve in and why?

A
  • like dissolved like

- polar solvents e.g. water

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32
Q

When an ionic solid dissolves in water, how do the ions interact with the water molecules?

A

+ve ion will be attracted to slight -ve (δ-) oxygen in water molecules

  • ve ions will be attracted to slightly +ve (δ+) hydrogens in water molecules
  • water molecules will completely surround the ions
  • enthalpy change occurs when new bonds form between ions and water molecules
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33
Q

Is hydration an endothermic or exothermic process?

A

exothermic

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34
Q

What factors affect the enthalpy change of hydration?

A
  • size of ions involved

- charges on the ions

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35
Q

How does the size of ions affect the enthalpy change of hydration?

A
  • smaller atomic radii so can get closer to water molecules
  • able to attract more strongly
  • on hydration more energy released so more exothermic
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36
Q

How does the charge on the ion affect the enthalpy change of hydration?

A
  • higher charge = more attraction with water molecule

- more -ve/exothermic

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37
Q

How can an ionic solid and gaseous ions be linked on a Born-Haber cycle?

A
  • lattice enthalpy

- standard enthalpy of solution along w standard enthalpy of hydration

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38
Q

What are the features of a Born-Haber cycle where enthalpy of solution, hydration and lattice enthalpy are represented?

A
  • ionic solid at bottom of cycle
  • gaseous ions at top of cycle
  • route via lattice enthalpy shown on left
  • route via enthalpies of solution and hydration shown on right
  • enthalpy change of solution (upwards arrow) doesn’t correspond to overall enthalpy change (downwards arrow): upwards = endothermic, downwards = exothermic
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39
Q

What does ‘entropy’ mean?

A

Measure of dispersal of energy in a system. The greater the entropy value, the more disordered the system
more spreading out of energy = higher entropy
more random arrangement of particles = higher entropy

40
Q

What does ‘standard entropy of a substance’ mean?

A

The entropy content of one mole of the substance under standard conditions

41
Q

What does ‘standard entropy change of reaction’ mean?

A

The entropy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactant and products being in their standard states

42
Q

What does a higher entropy value represent about a system?

A

more disordered particles in that system

43
Q

Order the 3 states in ascending entropy value.

A

solid < liquid < gas (gas is the most disordered)

44
Q

Why do all substances above 0K have a certain degree of disorder?

A

They are in constant motion

45
Q

What type of value is entropy?

A

+ve number above 0

46
Q

What is the entropy for perfect crystals at 0K?

A

0JK-1mol-1

47
Q

Describe the link between entropy and thermodynamic stability.

A
  • most substances thermodynamically stable at lowest energy state (aka lowest entropy)
  • entropy always tends to increase
  • more likely a disordered system will be found than an ordered one
48
Q

Give one example where entropy would be negative.

A

water freezing, as liquid –> solid it becomes more ordered w lower levels of energy dispersal

49
Q

How does temperature affect entropy?

A
  • entropy of pure substances increases with increasing temperature
  • values of entropy given per kelvin as entropy dependent on temperature
  • particles at higher temperature have higher energy and move more
  • arrangement of particles at higher temperatures becomes more random
  • entropy of solids < entropy of liquids < entropy of solids
  • ie when a liquid boils its entropy increases
50
Q

Describe the entropy changes when dissolving ionic solids.

A
  • if reaction results in products allowing more disorder (more ways for energy to be arranged/dispersed) entropy will increase
  • if solid ionic lattice dissolved, ions can spread out and positions of ions are far more disordered than within the lattice so entropy increases
51
Q

How does the number of gas molecules affect entropy?

A
  • increase in number of gas molecules causes an increase in entropy
  • decrease in number of gas molecules causes a decrease in entropy
52
Q

What calculation can be used to calculate the standard entropy change of reaction? What does a positive and negative result show?

A

ΔS=ΣS(standard)products - ΣS(standard)reactants

  • if change makes a system more random (entropy increases) ΔS is positive
  • if change makes a system more ordered (entropy decreases) ΔS is negative
53
Q

What does ‘free energy change, ΔG’ mean and what equation is used to represent this?

A

The balance between enthalpy, entropy and temperature for a process:
ΔG=ΔH-TΔS

54
Q

What value of ΔG is needed for a reaction to happen spontaneously?

A

ΔG<0

55
Q

What equation can be used to calculate the total change in entropy?

A

ΔS total = ΔS system + ΔS surroundings

all standard

56
Q

What conditions need to be met for a reaction to occur spontaneously?

A
  • ΔG<0

- total change in entropy must be positive

57
Q

What can make a reaction with a negative change in entropy happen spontaneously?

A

If entropy of surroundings is positive enough to make the total change in entropy positive

58
Q

What is used to measure free energy change? Using this, what is free energy dependent on?

A

-Gibbs’ equation
ΔG=ΔH-TΔS
-dependent on total entropy and enthalpy changes that occur, as well as on temperature

59
Q

How do large increases in entropy affect ΔG?

A

Decreases as ‘-TΔS’ is larger

60
Q

How do large negative values in ΔH affect ΔG?

A

More negative values of ΔG

61
Q

If ΔH is negative and ΔS is positive, what value of ΔG will you get? What does this show about the feasibility of a spontaneous reactions?

A
  • ΔG is always negative

- spontaneous reaction always feasible

62
Q

If ΔH is positive and ΔS is negative, what value of ΔG will you get? What does this show about the feasibility of a spontaneous reactions?

A
  • ΔG is always positive

- spontaneous reaction never feasible

63
Q

If ΔH is negative and ΔS is negative, what value of ΔG will you get? What does this show about the feasibility of a spontaneous reactions?

A
  • ΔG negative at low temperatures

- spontaneous reaction feasible at low temperatures

64
Q

If ΔH is positive and ΔS is positive, what value of ΔG will you get? What does this show about the feasibility of a spontaneous reactions?

A
  • ΔG negative at high temperates

- spontaneous reaction feasible at high temperatures

65
Q

Are exothermic reactions usually spontaneous or not and why?

A
  • generally spontaneous
  • negative ΔH value usually able to make ΔG negative, even is ΔS is positive (ΔH usually larger value, measured in kJ against ΔS’ J)
66
Q

Are endothermic reactions usually spontaneous or not and why?

A

Generally only spontaneous if entropy positive and temperature high enough to make TΔS large and positive (greater than ΔH)

67
Q

What are the limitations of using ΔG to predict feasibility?

A

-doesn’t take into account kinetic factors:
>reaction may have a high activation energy (energy needs to be initially supplied to overcome this i.e. igniting fuel as thermodynamically should happen spontaneously)
>rate of reaction may be extremely slow
-ΔG>0 reactions may take place however by changing the temperature

68
Q

What is a redox equation essentially made up of?

A
  • 2 half equations
  • one oxidation of a species
  • one reduction of a species
69
Q

What MUST be the same on both sides in a redox equation?

A

the electrons so they can cancel out

70
Q

Why do redox titrations not always need an indicator substance?

A

Many of the substances self indicate - they change colour between oxidation states

71
Q

How can manganate (VII) ions be self indicating in an redox titration?

A

MnO4 - ions are purple, reduced to Mn2+ (colourless)

72
Q

Describe the redox titration between Fe2+ and MnO4 -.

A

-manganate(VII) oxidising agent (from potassium permanganate(VII), KMnO4)
-usually occurs in present of H+ ions (preferably H2SO4 as HCl reacts w MnO4-: MnO4-(aq) + 8H+(aq) + 5e- –> Mn2+(aq) + 4H2O(l))
-manganate(VII) can oxidise substances other than iron(II) ions
MnO4-(aq) + 8H+ + 5Fe2+(aq) –> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
-end point when faint pink colour disappears (manganate added from burette) as all Fe2+ reacted so no more MnO4- can be reduced
-determine conc of iron (/other ions) in unknown solution or % composition of solid sample of compound/alloy

73
Q

Describe the redox titration between I2 and S2O3 2-.

A

-I2 has dark blue-black colour in presence of starch which goes when reduced to I- ions
2S2O3 2-(aq) + I2(aq) –> S4O6 2-(aq) + 2I-(aq)
-blue-black colour remains as long as theres iodine
-once all reacted, this will disappear (end point)
-can use aq I- ions and aq S2O3 2- to determine conc of unknown reducible species (often needs reaction between unknown oxidising agent and I- ions)

74
Q

What general calculations should be done when you get an unfamiliar redox system?

A
  • determine mol used of any substances of known conc and vol
  • identify reaction stoichiometry in balanced equations
  • decide amount that have reacted for known substances and deduce amounts of unknown solution
  • using equations calculate unknown quantities/concs
75
Q

How can redox titrations be used to estimate the copper content of solutions and alloys (e.g. brass, bronze)?

A

-any solid reacted w nitric acid to produce solution of Cu2+ ions, which reacts w KI(aq) Cu2+ ions mixed w I- ions and redox titration occurs
2I- –> I2 + 2e-
Cu2+ + e- –> Cu+
2Cu2+(aq) + 4I-(aq) –> 2CuI(s) + I2(aq)
-produces light brown/yellow solution and white ppt, CuI but ppt appears light brown from solution
-CuI/I2 mixture can be titrated against Na2S2O3 of known conc
-as I2 reacts, it gets paler during titration
-when colour is a pale straw colour, a small amount of starch is added to help identify the end point
-blue-black colour forms which disappears sharply at end point when all I2 reacted
-further calculations

76
Q

What is meant by standard electrode potential of a half cell?

A

The emf of a half cell compared with a standard hydrogen cell, measured at 298K w solution concs of 1moldm-3 and a gas pressure of 100kPa

77
Q

Describe the half cell of a metal.

A
  • metal electrode
  • metal ions in solution
  • forward reaction involves electron gain
  • reverse reaction involves electron loss
  • convention has electrons on left hand side
78
Q

Describe the hydrogen half cell.

A
  • platinum electrode (inert) - allows e- to pass into or out of half cell via a connecting wire
  • 2H+(aq) + 2e- ⇌ H2(g)
  • surface of Pt electrode coats w Pt black, a spongy coating in which electrons can be transferred between non metal and its ions
  • hydrochloric acid of 1 moldm-3 (source of H+)
  • hydrogen has, H2(g), at 100kPa
79
Q

Describe metal ion/metal ion half cells.

A
  • same element in different oxidation states
    e. g. Fe3+(aq) + e- ⇌ Fe2+(aq)
  • Pt electrode
  • solution w same conc of each metal ion (i.e. 1 moldm-3)
80
Q

What does a larger cell potential show?

A

More electrons ‘pushed around’ the cell

81
Q

How do you calculate the standard electrode potential of a half cell?

A
  • connect it to hydrogen half cell

- emf (electromotive force) measures tendency for different half cells to accept or release electrons

82
Q

What is the emf value of a hydrogen half cell?

A

0V - so can be used as a reference to other half cells

83
Q

What is the purpose of the wire in a cell?

A

Allows electrons carrying charge to flow through it

84
Q

What is the purpose of the salt bridge in a cell?

A
  • connects 2 solutions and allows ions carrying charge to be transferred between half cells
  • usually made from filter paper soaks in aqueous solution of an ionic substance, usually KNO3(aq)
85
Q

How can you work out which direction each half equation is going in a cell?

A

Greater E standard value (more +ve) = greater tendency to release electrons (go right)

86
Q

How is the electrode potential measured?

A

high resistance voltmeter (negligible current flows so pd can be measured)

87
Q

What equation do you use to calculate standard cell potential?

A

E standard (+ve terminal) - E standard (-ve terminal) = E standard (cell)

88
Q

What value of the standard cell potential would indicate that the reaction is feasible?

A

greater than 0V (reality 0.4V)

89
Q

What are the limitations of predicting cell potential from cell potentials?

A
  • not standard conditions

- slow reaction rate from high activation energy (so reaction won’t take place)

90
Q

How would increasing the copper ion concentration (i.e. not standard conditions) affect the electrode potential of the half cell?

A

Cu2+(aq) + 2e- –> Cu(s)

  • equilibrium opposes change by moving to right
  • e- removed from equilibrium
  • electrode potential becomes less negative (more positive)
91
Q

What is a non rechargeable cell?

A

Provides electrical energy until chemical have reacted to such an extent that the voltage falls. Is then ‘flat’ and discarded

92
Q

What is a rechargeable cell?

A
  • chemicals in cell react, providing electrical energy
  • cell reaction can be reversed during recharging
  • chemicals then regenerated and can be used again
93
Q

What are some common examples of rechargeable cells?

A
  • nickel-cadmium (Ni-Cad) batteries used in rechargeable batteries
  • lithium-ion and lithium, polymer batteries, used in laptops
94
Q

What are the 3 main types of cell?

A
  • non rechargeable cells
  • rechargeable cells
  • fuel cells
95
Q

What is a fuel cell?

A
  • cell reaction uses external supplies of fuel and an oxidant, which are consumed and need to be continuously supplied
  • cell continues to provide electrical energy so long as theres a supply of fuel and oxidant
96
Q

What are the benefits and drawbacks of lithium cells?

A

+high levels of battery life
-toxicity on being ingested
-rapid discharge of current, which can cause fires or explosions
(hence restriction on transport of lithium based batteries and limited sales of them to individual consumers in some countries)

97
Q

How do hydrogen based modern fuel cells work?

A

-use reaction w oxygen to create a voltage
-reactants flow in and products flow out while electrolyte remains in cell
-can operate virtually continuously if fuel and oxygen continue to flow in
-don’t need to be recharged
2H2O(l) + 2e- ⇌ H2(g) + 2OH-(aq)
1/2O2(g) + H2O(l) + 2e- ⇌ 2OH-
H2(g) + 1/2O2(g) –> H2O(l)

Chemistry Year 12/13 (18 decks)

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