Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

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

Module 5.2 - Energy Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

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

A
  • exothermic

- bonds are made

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

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

A

K(s) + 1/2Cl2(g) –> KCl(s)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

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

A

Endothermic as bonds have to be broken

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What does ‘standard enthalpy change of atomisation’ mean?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

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

A

K(s) –> K(g)

1/2Cl2(g) –> Cl(g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What does ‘first electron affinity’ mean?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Give the equation for the first electron affinity of chlorine.

A

Cl(g) + e- –> Cl-(g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

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

A

sum of clockwise enthalpies = sum of anticlockwise enthalpies

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Is enthalpy change of solution endothermic or exothermic?

A

Can be either

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

How does the charge of an ion affect lattice enthalpy?

A
  • higher charges causes greater electrostatic attraction

- more exothermic (more -ve)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

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

A

very small, highly charged

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

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

A
  • like dissolved like

- polar solvents e.g. water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

Is hydration an endothermic or exothermic process?

A

exothermic

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

What factors affect the enthalpy change of hydration?

A
  • size of ions involved

- charges on the ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions