Ch5.2 Energy Flashcards
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
2
Q
Describe lattice enthalpy value
A
- theoretical value
- indicates relative strength of ionic bonds
3
Q
What does a more exothermic lattice enthalpy indicate?
A
- stronger ionic bonds
- stronger electrostatic attractions
- higher melting/boiling points
4
Q
When do the most exothermic lattice enthalpies arise?
A
- when ions are small and have large charges
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
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
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
8
Q
Define second ionisation energy
A
- one mole of gaseous 2+ ions is formed from one mole of gaseous 1+ ions
- endothermic
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
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
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
12
Q
Define standard enthalpy change of hydration
A
the enthalpy change when one mole of gaseous ions are dissolved in water
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
14
Q
Which has higher entropy, gases or solids?
A
gases
15
Q
What does entropy always tend to do?
A
- increase
- eg water always tends to evaporate
- eg gases spread to fill a container
16
Q
What is the standard entropy of a substance?
A
the entropy content of one mole of a substance under standard conditions
17
Q
What are the units for entropy?
A
J K-1 mol-1
18
Q
What are standard conditions?
A
298K 100kPa
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
20
Q
What effect does temp have on entropy?
A
- the entropy of pure substances increases with temp
21
Q
Describe how entropy changes when ionic solids are dissolved
A
- the ions spread out and become more disordered
- entropy increases
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
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
24
Q
Formula for change in entropy
A
ΔS = ΣS° (products) - ΣS° (reactants)
25
What are spontaneous reactions?
- 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
26
Equation for total change in entropy
ΔS° (total) = ΔS° (system) + ΔS° (surroundings)
27
What must be true for a spontaneous reaction to occur?
- 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
28
What is the free energy change?
ΔG, is the balance between enthalpy, entropy and temp
29
What is the equation for Gibbs' free energy?
ΔG = ΔH - TΔS
30
What must ΔG be for a reaction to occur spontaneously?
ΔG < 0
31
Which type of reaction is generally spontaneous?
- exothermic
- the negative value of ΔH is usually still able to make ΔG negative even if entropy is positive
32
Which type of reactions are likely not to be spontaneous?
- endothermic
- only spontaneous if the entropy is positive and the temp is high enough to make TΔS greater than ΔH
33
Will all reactions with negative ΔG spontaneously react?
- no
- it also depends on the size of the activation energy and the rate of reaction
34
What does a positive ΔG mean?
- the reaction is not feasible too be spontaneous but can be made feasible by changing the temp of the reaction
35
What are redox titrations used for?
- 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
What colour is MnO4- before being reduced?
purple
37
What colour is MnO4- after being reduced from +7 to +2?
colourless (very pale pink)
38
What is MnO4- usually reduced in the presence of?
H+ ions (usually from H2SO4)
39
What is MnO4- usually used to oxidise?
Fe2+ ions to Fe3+ ions
40
Describe what happens in a redox titration
(specificto manganate VII and iron II)
- 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
What can iodine be titrated against?
- thiosulphate ions
- forms iodide ions and tetrathionate ions
42
Explain the colour change in a titration involving thiosulphate ions and iodine
- 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
Describe the method used to find the concentration of a solution of an oxidising agent that will oxidise iodide ions to iodine
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
What do batteries and cells rely on?
redox reactions
45
Describe a half cell
- 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
What are the 2 methods for making a half cell?
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
Describe a copper half cell
- 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
Describe a hydrogen half cell
- hydrogen gas in contact with a solution of its aqueous ions (H+)
- an equilibrium is present between the hydrogen gas and H+ ions
49
Explain when a platinum electrode is used in a hydrogen half cell
- 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
Describe an Fe2+/Fe3+ half cell
- solution containing 2+ and 3+ ions
- they must be of equal concentrations (EQUIMOLAR)
- an equilibrium is present between the ions
51
What does the electrode potential of a half cell tell us?
its tendency to lose or gain electrons in the equilibrium
52
Define standard electrode potential
- 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
How do you work out the standard electrode potential of a half cell?
- 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
Why is the standard hydrogen electrode significant?
- it is used as a reference point for all other half cells to be measured against
- it has an emf of 0V
55
Describe the method of joining 2 half cells to measure standard electrode potentials
- 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
What is a salt bridge?
- 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
What do negative and positive values for electrode potential indicate?
- 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
Which electrode potential will be the forward reaction?
the more positive one
59
What does a larger electrode potential cell value inidcate?
the further the half equations move from equilibrium
60
State some limitations to predicting feasibility of reactions
- 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
How can you make an electrochemical cell with a large voltage?
- you need 2 half cells with different standard electrode potentials
- the larger the difference the larger the emf of the cell
62
Describe non-rechargeable cells
- provide electrical energy until all of the chemicals have reacted
- it is then said to be flat
63
Describe rechargeable cells
- provide electrical energy when the chemicals react
- chemical reaction can be reversed during recharging so the chemicals can react again
64
Give 2 examples of rechargeable cells
- nickel-cadmium battery in rechargeable batteries
- lithium ion/polymer batteries in laptops
65
State the benefits and risks of lithium batteries
benefits
- high level of battery life
- less waste
risks
- toxic if ingested
- can cause fires and explosions if current is discharged too quickly
66
Describe fuel cells
- 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
Describe how the hydrogen-oxygen fuel cell works
- 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
State advantages of the hydrogen-oxygen fuel cell
- water is the only waste product
- high efficiency
69
State disadvantages of the hydrogen-oxygen fuel cell
- expensive to produce
- hydrogen is highly flammable
- hydrogen is difficult to store
- hydrogen is difficult to manufacture
