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Chemistry > Ch5.2 Energy > Flashcards

Ch5.2 Energy Flashcards

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1
Q

Define 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

Describe lattice enthalpy value

A
  • theoretical value
  • indicates relative strength of ionic bonds
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3
Q

What does a more exothermic lattice enthalpy indicate?

A
  • stronger ionic bonds
  • stronger electrostatic attractions
  • higher melting/boiling points
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4
Q

When do the most exothermic lattice enthalpies arise?

A
  • when ions are small and have large charges
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5
Q

Define standard enthalpy change of formation
exo or endo

A
  • one mole of a compound is formed from its constituent elements in their standard states
  • always exothermic
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6
Q

Define standard enthalpy change of atomisation
exo or endo

A
  • one mole of gaseous atoms is formed from its element in its standard state
  • always endothermic
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7
Q

Define first ionisation energy
exo or endo

A
  • one mole of gaseous 1+ ions is formed from one mole of gaseous atoms
  • endothermic
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8
Q

Define second ionisation energy

A
  • one mole of gaseous 2+ ions is formed from one mole of gaseous 1+ ions
  • endothermic
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9
Q

Define first electron affinity

A
  • opposite of ionisation energy
  • one mole of gaseous 1- ions is formed from 1 mole of gaseous atoms
  • exothermic
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10
Q

Define second electron affinity
endo or exo

A
  • one mole of gaseous 2- ions is formed from one mole of gaseous 1- ions
  • endothermic because the electron is repelled by the 1- ion
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11
Q

Define standard enthalpy change of solution

A

the enthalpy change when one mole of a solute is completely dissolved in water under standard conditions

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12
Q

Define standard enthalpy change of hydration

A

the enthalpy change when one mole of gaseous ions are dissolved in water

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13
Q

What is entropy?

A
  • entropy (S) is the quantitative measure of the degree of disorder in a system
  • energy disperses as the system becomes more disordered
  • highly ordered=low entropy
  • highly disordered=high entropy
  • as disorder increases, entropy increases
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14
Q

Which has higher entropy, gases or solids?

A

gases

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15
Q

What does entropy always tend to do?

A
  • increase
  • eg water always tends to evaporate
  • eg gases spread to fill a container
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16
Q

What is the standard entropy of a substance?

A

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

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17
Q

What are the units for entropy?

A

J K-1 mol-1

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18
Q

What are standard conditions?

A

298K 100kPa

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19
Q

Is entropy a positive or negative value?

A
  • entropy is always positive when substances are above 0K
    as temp increases, disorder increases and energy soreads out so entropy increases
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20
Q

What effect does temp have on entropy?

A
  • the entropy of pure substances increases with temp
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21
Q

Describe how entropy changes when ionic solids are dissolved

A
  • the ions spread out and become more disordered
  • entropy increases
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22
Q

Describe how the number of gas molecules in a reaction affects entropy

A
  • if the number of gas molecules changes during a reaction, entropy changes
  • if the number of gas molecules increases after a reaction, the entropy increases and vice versa
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23
Q

Define standard entropy change of reaction

A
  • the entropy change when a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products being in their standard states
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24
Q

Formula for change in entropy

A

ΔS = ΣS° (products) - ΣS° (reactants)

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25
Q

What are spontaneous reactions?

A
  • occur instantaneously
  • entropy always increases
  • there are spontaneous reactions which appear to lower entropy
  • to explain this, the total entropy change of the system and the surroundings need to be considered
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26
Q

Equation for total change in entropy

A

ΔS° (total) = ΔS° (system) + ΔS° (surroundings)

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27
Q

What must be true for a spontaneous reaction to occur?

A
  • ΔS° (total) must be positive
  • in reactions where the system entropy change decreases, if the surrounding entropy change increases enough the reaction can still be spontaneous
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28
Q

What is the free energy change?

A

ΔG, is the balance between enthalpy, entropy and temp

29
Q

What is the equation for Gibbs’ free energy?

A

ΔG = ΔH - TΔS

30
Q

What must ΔG be for a reaction to occur spontaneously?

A

ΔG < 0

31
Q

Which type of reaction is generally spontaneous?

A
  • exothermic
  • the negative value of ΔH is usually still able to make ΔG negative even if entropy is positive
32
Q

Which type of reactions are likely not to be spontaneous?

A
  • endothermic
  • only spontaneous if the entropy is positive and the temp is high enough to make TΔS greater than ΔH
33
Q

Will all reactions with negative ΔG spontaneously react?

A
  • no
  • it also depends on the size of the activation energy and the rate of reaction
34
Q

What does a positive ΔG mean?

A
  • the reaction is not feasible too be spontaneous but can be made feasible by changing the temp of the reaction
35
Q

What are redox titrations used for?

A
  • to determine the amount of a species which is oxidised or reduced in a reaction
  • the reducing or oxidising agent that you know the conc of is put in the burette and titrated against a chemical which is being oxidised or reduced
36
Q

What colour is MnO4- before being reduced?

A

purple

37
Q

What colour is MnO4- after being reduced from +7 to +2?

A

colourless (very pale pink)

38
Q

What is MnO4- usually reduced in the presence of?

A

H+ ions (usually from H2SO4)

39
Q

What is MnO4- usually used to oxidise?

A

Fe2+ ions to Fe3+ ions

40
Q

Describe what happens in a redox titration
(specificto manganate VII and iron II)

A
  • acidified manganate (VII) is added to the iron (II) until all the Fe2+ is oxidised
  • a very faint pink colour will then appear because of unreacted manganate
41
Q

What can iodine be titrated against?

A
  • thiosulphate ions
  • forms iodide ions and tetrathionate ions
42
Q

Explain the colour change in a titration involving thiosulphate ions and iodine

A
  • starch is used to more easily identify the endpoint as the natural colour change can be hard to see if there are other compounds present
  • when there is iodine still present, it is blue-black
  • when all the iodine has reacted the blue-black will disappear and the solution should be colourless but it may appear white or cream if there’s other compounds present
43
Q

Describe the method used to find the concentration of a solution of an oxidising agent that will oxidise iodide ions to iodine

A
  1. react iodide ions with the oxidising agent to form iodine
  2. titrate the iodine against a known conc of sodium thiosulphate
  3. calculate the moles of iodine that must have reacted with the sodium thiosulphate
  4. calculate the conc of the oxidising agent
44
Q

What do batteries and cells rely on?

A

redox reactions

45
Q

Describe a half cell

A
  • when you look at the redox reactions that are part of a cell each half equation involved is called a half cell
  • half cells can be combined to make a cell
  • a half cell contains an element in 2 different oxidation states
46
Q

What are the 2 methods for making a half cell?

A
  1. putting a metal or non-metal in contact with their ions that are in aqueous solution
  2. having ions of the same element in different oxidation states within an aqueous solution
47
Q

Describe a copper half cell

A
  • copper metal in an aqueous solution of Cu2+ ions
  • an equilibrium is present between the copper metal and the Cu2+ ions at the copper metal’s surface
48
Q

Describe a hydrogen half cell

A
  • hydrogen gas in contact with a solution of its aqueous ions (H+)
  • an equilibrium is present between the hydrogen gas and H+ ions
49
Q

Explain when a platinum electrode is used in a hydrogen half cell

A
  • used if you want to connect the half cell to another half cell
  • platinum is used as it’s inert and will only transfer electrons to and from the solution
  • the surface of the electrode is coated with ‘platinum-black’ to help electrons to be transferred between the non-metal and its ions
50
Q

Describe an Fe2+/Fe3+ half cell

A
  • solution containing 2+ and 3+ ions
  • they must be of equal concentrations (EQUIMOLAR)
  • an equilibrium is present between the ions
51
Q

What does the electrode potential of a half cell tell us?

A

its tendency to lose or gain electrons in the equilibrium

52
Q

Define standard electrode potential

A
  • the e.m.f. of a half cell compared with a
    standard hydrogen half cell. This is measured under standard conditions (1 mol dm-3 solution
    concentrations, 298K and 1 atm)
53
Q

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

A
  • you must connect it to a hydrogen half cell
  • one half cell will then release electrons while the other will gain them
  • the difference in the electrode potential of the 2 half cells is measured using a voltmeter
54
Q

Why is the standard hydrogen electrode significant?

A
  • it is used as a reference point for all other half cells to be measured against
  • it has an emf of 0V
55
Q

Describe the method of joining 2 half cells to measure standard electrode potentials

A
  • wire connects 2 half cells and allows electrons to be transferred between them
  • salt bridge allows ions to be transferred between half cells
  • both of these are needed to complete the circuit otherwise the current wouldn’t flow
56
Q

What is a salt bridge?

A
  • filter paper soaked in an aqueous solution of an ionic compound that doesn’t react with either half cell solution
  • eg. KNO3 or NH4NO3
57
Q

What do negative and positive values for electrode potential indicate?

A
  • negative means the backward reaction occurs in comparison to the hydrogen electrode and it has a greater tendency to oxidation
  • positive means reduction occurs in comparison to the hydrogen electrode and it has a greater tendency to reduction
58
Q

Which electrode potential will be the forward reaction?

A

the more positive one

59
Q

What does a larger electrode potential cell value inidcate?

A

the further the half equations move from equilibrium

60
Q

State some limitations to predicting feasibility of reactions

A
  • electrode potentials are for standard conditions only with solutions being 1mol/dm3
  • if concs were changed, the position of equilibrium would change due to le chateliers principle so electrons would be removed/added and electrode potential would change
61
Q

How can you make an electrochemical cell with a large voltage?

A
  • you need 2 half cells with different standard electrode potentials
  • the larger the difference the larger the emf of the cell
62
Q

Describe non-rechargeable cells

A
  • provide electrical energy until all of the chemicals have reacted
  • it is then said to be flat
63
Q

Describe rechargeable cells

A
  • provide electrical energy when the chemicals react
  • chemical reaction can be reversed during recharging so the chemicals can react again
64
Q

Give 2 examples of rechargeable cells

A
  • nickel-cadmium battery in rechargeable batteries
  • lithium ion/polymer batteries in laptops
65
Q

State the benefits and risks of lithium batteries

A

benefits
- high level of battery life
- less waste

risks
- toxic if ingested
- can cause fires and explosions if current is discharged too quickly

66
Q

Describe fuel cells

A
  • cell reaction uses an external supply of a fuel and oxidant
  • the cell will continue to provide electrical energy as long as fuel and oxidant don’t run out
  • fuel is usually based on hydrogen or hydrogen-rich fuels eg. methanol
67
Q

Describe how the hydrogen-oxygen fuel cell works

A
  • hydrogen and oxygen flow into the cell and the water flows out of the cell while the electrolyte stays inside (KOH)
  • the hydrogen reacts with the oxygen to create a voltage
68
Q

State advantages of the hydrogen-oxygen fuel cell

A
  • water is the only waste product
  • high efficiency
69
Q

State disadvantages of the hydrogen-oxygen fuel cell

A
  • expensive to produce
  • hydrogen is highly flammable
  • hydrogen is difficult to store
  • hydrogen is difficult to manufacture

Chemistry (17 decks)

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