Module 5.2 Flashcards
1
Q
Define the term standard lattice enthalpy
A
The enthalpy change when one mole of ionic lattice is formed from its gaseous ions under standard conditions
2
Q
Give an example reaction equation that shows standard lattice enthalpy
A
Sodium ion(g) + Chloride ion(g) → NaCl(s)
Calcium ion(g) + Oxygen ion (g) → CaO(s)
3
Q
Define the term standard enthalpy change of formation
A
The enthalpy change when one mole of compound is formed from its elements in their defined standard states under standard conditions
4
Q
Give an example reaction equation that shows standard enthalpy change of formation
A
K(s) + 1/2F2(g) → KF(s)
2C(s) + 3H2(g) + 1/2O2 (g) → C2H5OH
5
Q
Define the term standard enthalpy change of atomisation
A
The enthalpy change when one mole of gaseous atoms is from from its elements in their defined standard states under standard conditions
6
Q
Give an example reaction equation that shows standard enthalpy change of atomisation
A
K(s) → K(g)
1/2Br2 (s) → Br2(g)
7
Q
Define the term first ionisation energy
A
The amount of energy required to remove one mole of electrons from one mole of gaseous atoms.
8
Q
Give an example reaction equation that shows first ionisation energy
A
Na → Na^+ + e^-
O → O^+ + e^-
9
Q
Define the term second ionisation energy
A
The amount of energy that accompanies the formation of one mole of gaseous 2+ ions from 1 mole of gaseous 1+ iions
10
Q
Give an example reaction equation that shows second ionisation energy
A
Na^+ → Na^2+ + e^-
O^+ → O^2+ + e^-
11
Q
Define the term first electron affinity
A
The enthalpy change that accompanies the formation of one mole of gaseous 1^- ions from gaseous atoms
12
Q
Give an example reaction equation that shows first electron affinity
A
Cl(g) + e^- → Cl^-(g)
1/2F2(g) + e^- → F^-(g)
13
Q
Define the term second electron affinity
A
The enthalpy change that accompanies the formation of one mole of gaseous 2^- ions from one mole of gaseous 1^- ions.
14
Q
Give an example reaction equation that shows first electron affinity
A
Cl^-(g+ e^- → Cl^2-(g)
F^-(g) + e^- → F^-(g)
15
Q
What are the standard conditions
A
100kPa and 298Kelvin
16
Q
Is Lattice Enthalpy always endothermic or always exothermic?
A
Always EXOTHERMIC
17
Q
What factors determine lattice enthalpy?
A
Ionic size and Ionic charge
18
Q
How does a decreasing ionic size affect lattice enthalpy?
A
Smaller ions have a higher charge density. This means that the ions attract to each other more strongly. Therefore, more energy is released when bonds form. Therefore, the smaller the ionic radius, the bigger the lattice enthalpy
19
Q
How does an increasing ionic size affect lattice enthalpy?
A
Bigger ions have a smaller charge density. This means that the ions attract to each other more weakly. Therefore, less energy is released when bonds form. Therefore, the bigger the ionic radius, the smaller the lattice enthalpy
20
Q
How does an increasing ionic chargevaffect lattice enthalpy?
A
The higher the charge on the ions, the more energy is released when an ionic lattice forms. This is due to the stronger electrostatic forces between the ions. Therefore, the higher the ionic charge, the larger the lattice enthalpy
21
Q
How does a decreasing ionic charge affect lattice enthalpy?
A
The smaller the charge on the ions, the less energy is released when an ionic lattice forms. This is due to the weaker electrostatic forces between the ions. Therefore, the smaller the ionic charge, the smaller the lattice enthalpy
22
Q
Higher lattice enthalpy means a higher or lower negative value?
A
Higher
23
Q
What factors affect ionisation energy?
A
Nuclear charge
Shielding
Atomic radius
24
Q
Greater nuclear charge = _______ ionisation energy. Why?
A
Higher. This is because the outermost electrons are more strongly attached to the nucleus. Therefore, more energy is needed to remove one mole of electrons from a gaseous atom.
25
Greater shielding = _______ ionisation energy. Why?
Lower. This is because the outermost electrons are less strongly attached to the nucleus. Therefore, less energy is needed to remove one mole of electrons from a gaseous atom.
26
Greater atomic radius = _______ ionisation energy. Why?
Lower. This is because the outermost electrons are further away from the nucleus. Therefore, they are less strongly attached to the nucleus. This means that less energy is needed to remove one mole of electrons from a gaseous atom.
27
What is a Born-Haber Cycle and what is it used to do?
A Born-Haber cycle is an indirect cycle which allows us to find lattice enthalpy for a reaction.
28
What does a generic Born-Haber cycle look like?
TOP
First IE
↑ Electron affinity
Atomisation ↓
↑ ↓
Atomisation ↓
↓ ↓
Formation - Formation - Formation - Formation
BOTTOM
29
What thing must you include in a Born-Haber cycle?
State symbols
Direction of arrows
Diatomic?
Name of steps (e.g. enthalpy change of atomisation)
What do you times by 2 or 3 etc
- Exothermic or endothermic (Watch out for 2nd electron affinity)
- Balancing charges
30
In the Born-Haber cycle for a group 1 and group 7 element, do we times anything by 2 or 3?
NO
31
In the Born-Haber cycle for a group 2 and group 6 element, do we times anything by 2 or 3? If yes, what?
NO
32
In the Born-Haber cycle for a group 3 and group 6 element, do we times anything by 2 or 3? If yes, what?
Times by 2 : First IE, 2nd IE and 3rd IE (as there are 2 moles of Group 3 element)
Times by 3: First Electron Affinity and Second electron affinity (as there are 3 moles of Group 4 element)
33
In the Born-Haber cycle for a group 2 and group 7 element, do we times anything by 2 or 3? If yes, what?
Times by 2: Atomisation of the halogen and the first electron affinity (as there are 2 moles of group 7 element by 1 mole group 2 element)
34
Give the formula for Lattice Enthalpy
LE =
- Electron affinity - First IE - Enthalpy change of atomisation of element X - Enthalpy change of atomisation of element Y + Enthalpy change of formation
35
Define enthalpy change of solution
The enthalpy change when one mole of a solute dissolves in water under standard conditions
36
Give 3 example reactions equations that show the enthalpy change of solution
NaCl(s) → Na^+(g) + Cl^-(g)
MgF2l(s) → Mg^2+(g) + F^-(g)
MgCO3(s) → Mg^2+(g) + CO3^2-(g)
37
What formula can we use to find the enthalpy change of solution?
Enthalpy change of solution =
Enthalpy change of hydration of X + Enthalpy change of hydration of Y - Standard lattice enthalpy
*Be careful with signs, take time. This is easy marks.
38
Define enthalpy change of hydration
The enthalpy change when one mole of gaseous ions is dissolved in water under standard conditions
39
Give 3 example reactions equations that show the enthalpy change of hydration
Na^+(g) → Na^+(aq)
K^+(g) → K^+(aq)
O^2-(g) → O^2-(aq)
40
Is enthalpy change of hydration, exothermic or endothermic?
Exothermic because bonds are being made with water
41
Is enthalpy change of solution, exothermic or endothermic?
It can be either
42
What happens when an ionic solid dissolves in water?
When an ionic solid dissolves in water, the bonds between the ions BREAK to give gaseous ions. This is ENDOTHERMIC. Furthermore, bonds between the gaseous ions and the WATER are made (forms aqueous ions). This is EXOTHERMIC.
43
Give the Born- Haber cycle for a hydration/solution containing cycle
TOP
Gaseous atoms - Gaseous atoms - Gaseous atoms
↓ ↓
↓ ↓ Enthalpy change of hydration for X
↓ ↓
↓ ↓ Enthalpy change of hydration for Y
↓ ↓
↓ ↓ Enthalpy change of solution
↓ ↓
Ionic solid - Ionic solid - Ionic solid - Ionic solid
BOTTOM
44
How can you calculate lattice enthalpy from a Born-Haber cycle that consists of enthalpy change of hydration and enthalpy change of formation
LE + Solution = Hydration of X + Hydration of Y
Therefore LE = Hydration of X + Hydration of Y - Solution
45
If in the Born-Haber cycle that consists of enthalpy change of hydration and enthalpy change of solution, there are 2 moles of something, what do you do to its enthalpy change of hydration?
Times that by 2
46
What factors affect the enthalpy change of hydration?
Ionic size and Ionic charge
47
How does a decreasing ionic size affect affect enthalpy change of hydration?
Smaller ions have a higher charge density than bigger ions. They ATTRACT the water molecules BETTER and have a more exothermic enthalpy change of hydration. Therefore, the smaller the ion, the larger its enthalpy change of hydration (due to having a higher charge density and better water attraction
48
How does an increasing ionic size affect affect enthalpy change of hydration?
Bigger ions have a smaller charge density than smaller ions. They ATTRACT the water molecules WEAKER and have a less exothermic enthalpy change of hydration. Therefore, the bigger the ion, the smaller its enthalpy change of hydration (due to having a smaller charge density and weaker water attraction
49
How does an increasing ionic charge affect affect enthalpy change of hydration?
Ions with a higher ionic charge are better at ATTRACTING water molecules than those with lower charges. This means that the electrostatic attraction between the ion and the water molecule is STRONGER. This means that more energy is released when bonds are made, giving them a more EXOTHERMIC enthalpy change of hydration. Therefore, the higher the ionic charge, the higher the enthalpy change of hydration
50
How does a decreasing ionic charge affect affect enthalpy change of hydration?
Ions with a lower ionic charge are bad at ATTRACTING water molecules than those with higher charges. This means that the electrostatic attraction between the ion and the water molecule is WEAKER . This means that less energy is released when bonds are made, giving them a less EXOTHERMIC enthalpy change of hydration. Therefore, the lower the ionic charge, the lower the enthalpy change of hydration
51
Define the term entropy
A measure of dispersal in a system, which is greater, the more disordered the system
52
Higher entropy means?
More disorder (likely to be a gas)
53
Lower entropy means?
Less disorder (likely to be a solid)
54
How do we look at entropy?
In terms of the number of ways that energy can be shared and the number of ways that particles can be arranged.
55
More particles mean?
More entropy. This is because, the more particles we have, the more ways they and their energy can be arranged.
56
High entropy =
More disordered = Positive value = Likely to be a gas
57
Low entropy =
Less disordered = Negative value = Likely to be a solid
58
How is entropy affected by temperature?
An increase in temperature means the particles have more kinetic energy and so move more. This means that they become more disordered so entropy increases with temperature (Directly proportional relationship)
59
How does entropy change in terms of physical states?
As a substance changes from SOLID to LIQUID to GAS, it's entropy will INCREASE each time. This is because particles are becoming more spread out, therefore more disordered.
60
How does entropy change in terms of dissolving ionic lattices?
As an ionic lattice dissolves (becomes aqueous), the entropy will increase. This is because ions can spread out and the positions of the ions are far more disordered than within the lattice. This means that entropy INCREASES.
61
How does entropy change in terms of number of gaseous moles? (Give the 2 rules)
Rule 1) When a reaction occurs, if there are more moles of gas in the PRODUCTS, the entropy increases.
Rule 2)When a reaction occurs, if there are less moles of gas in the PRODUCTS, the entropy decreases.
62
Give the formula for change in Entropy
Change in entropy = Entropy of Products - Entropy of Reactants
63
When finding entropy change, what must we take into account?
Mole number
64
If change in entropy is positive what does it mean?
The system has become more disordered
65
If change in entropy is negative, what does it mean?
The system has become less disordered
66
What is the Gibbs Free Energy equation?
ΔG = ΔH - TΔS
Free energy change = Enthalpy change - (Temperature x Entropy change)
67
In the free energy change formula, what does everything stand for?
```
ΔG = Free energy change (J/mol)
ΔH = Enthalpy change (kJ/mol)
T = Temperature (Kelvin)
ΔS = Entropy change
```
68
What must we divide by 1000, before inputting it, into our free energy change formula?
ΔS
| converts it from J to kJ/mol
69
How do we find ΔH (Enthalpy change) ?
ΔH = Σ(enthalpy of products) - Σ(enthalpy of reactants)
70
How do we find ΔS (Entropy change)
ΔS = Σ(entropy of products) - Σ(entropy of reactants)
71
How do we find temperature using the Free energy equation?
T = ΔH / ΔS
Temperature = Enthalpy change / Entropy change
72
If ΔG is less than 0 (ΔG<0), what does it mean?
So ΔH is negative and ΔS is positive.
It means the reaction is feasible
73
If ΔG is more than 0 (ΔG>0), what does it mean?
So ΔH is positive, and ΔS is negative
It means that the reaction WILL NOT proceed
74
If ΔH is negative and ΔS is also negative, what does it mean?
The reaction is feasible when TΔS < ΔH.
The reaction is feasible at lower temperatures
75
If ΔH is positive and ΔS is also positive, what does it mean?
The reaction is feasible when TΔS > ΔH.
The reaction is feasible at high temperatures
76
If the reaction is feasible does it mean that will take place? Give 1 reason to support your answer
No it will not take place just because it's feasible.
| One reason for this could be unfavourable kinetics e.g. activation energy may be too high.
77
Describe the limitations of predictions made by ΔG about feasibility
Give 2 limitations.
The reaction may have a high activation energy. Therefore, even if the reaction is feasible, due to such a high Ea, the reaction may not even take place.
Standard Gibbs free energies are calculated for a specific set of conditions. It may be that the reaction proceeds normal under OTHER STANDARD CONDITIONS and perhaps not the current standard conditions.
78
What is meant by an oxidising agent?
A species that accepts electrons and gets reduced itself
79
Give 2 examples of an oxidising agent
Acidified potassium dichromate
| Potassium permanganate
80
What is meant by a reducing agent?
A species that donates electrons and get oxidised itself.
81
Give an example of a reducing agent
Sodium borohydride
82
What is a redox reaction?
A reaction where one species is oxidised and another is reduced.
83
What is a redox half equation?
An equation showing a species being either oxidised or reduced
84
What can we use to balance redox half equations?
H20
e^-
H^+ / OH^-
85
How do you construct redox equations using half equations?
- Balance out charges on both sides for each half equation (half equations don't need the same overall charge on each side but they overall equation does)
- Use H20 / e^- / H^+ / OH^- to balance the charges on both sides
86
In terms of a redox half equation, if the conditions are acidic what can we use?
H^+
87
In terms of a redox half equation, if the conditions are alkaline/basic what can we use?
OH^-
88
Common redox reaction half equation :
MnO4^- → 4H2O
MnO4^- + 8H^+ + 5e^- → Mn2+ + 4H2O
89
Finish off the common redox reaction half equation
Cr2O7^2- → Cr^3+
Cr2O7^2- + 14H^+ + 6e^- → 2Cr^3+ + 7H2O
90
Finish off the common redox reaction half equation
Fe^2+ → Fe^3+
Fe^2+ → Fe^3+ + e^-
91
Finish off the redox reaction half equation
Sn^2+ → Sn^4+
Sn^2+ → Sn^4+ + 2e^-
92
Finish off the redox reaction half equation :
SO2 → SO4^2-
SO2 +2H2O → SO4^2- + 4H^+ + 2e^-
93
Finish off the redox reaction half equation
VO2^+ → VO^2+
VO2^+ + 2H^+ + e^- → VO^2+ + H2O
94
In terms of electrons, how do we know an element has been oxidised?
If it LOSES electrons
95
In terms of electrons, how do we know an element has been reduced?
If it GAINS electrons
96
To get an our overall redox equation, list a few things we must know/do
- Multiply half equations to allow electrons to cancel out
- Same side groups = ADD
- Opposite side groups = SUBTRACT
- Check overall charge on both sides at the end
97
Give the Fe^2+/MnO4^- titration method
1) Using a pipette, measure out a quantity of the reducing agent (Fe^2+)
2) Add some dilute H2SO4 to the flask
3) Gradually add the oxidising agent (MnO4^-) to the reducing agent using a burette, swirling the conical flask around as you do so
4) Stop when the mixture in the flask just becomes TAINTED PINK. This point is known as the end point
5) Record the volume of the oxidising agent added. This gives us our titre.
6) Run a few extra titres and calculate the mean value of oxidising agent added using CONCORDANT titres.
98
Fe^2+/MnO4^- titration common exam question 1:
Why is ammonium iron sulfate suitable as a primary standard?
- Stable
| - Available in a highly pure form
99
Fe^2+/MnO4^- titration common exam question 2:
Why is sulfuric acid added to the iron(Ii) solution prior titration?
Acidic conditions are necessary for the fully reduction of the oxidising agent
100
Fe^2+/MnO4^- titration common exam question 3:
Rather than using H2SO4, could you use HNO3 instead? Explain your answer
We can't use nitric acid instead because it is a power oxidising agent.
101
Fe^2+/MnO4^- titration common exam question 4:
Rather than using H2SO4, could you use HCl instead? Explain your answer
We can't use HCl instead because chlorine gas will be produced when it reacts with MnO4^-.
Chlorine gas is toxic.
102
Fe^2+/MnO4^- titration common exam question 5:
In preparation for the titration, explain why the pipette and burette are rinsed with deionised water followed by a little of the solutions they contain
Deionised water removes any residual solutions in the burette and pipette.
The 2nd step is to remove any residual deionised water, and so avoid the dilution of the solutions, why they are added to the burette and pipette.
103
Fe^2+/MnO4^- titration common exam question 6
During the titration, the sides of the conicial flask are washed down with deionised water, why?
To ensure that all the MnO4^- added from the burette reacted with all the Fe^2+
104
Fe^2+/MnO4^- titration common exam question 7
What was the catalyst in the reaction?
Mn^2+ ions
105
Fe^2+/MnO4^- titration common exam question 8
Why was H2SO4 added in the making up of the ammonium iron (II) sulfate solution?
Without using H2SO4, ammonium iron sulfate would've been oxidised rather than the iron ions.
106
Give the I2/S2O3^2- titration method
1) Measure out a certain value of potassium iodide solution (KIO3)
2) Add this to an excess of acidified potassium iodide solution (KI)
3) Using a burette, add sodium thiosulfate to the flask, drop wise.
4) When the iodine colour fades from a pale yellow, add 2.00cm^3 of starch. ( the colour of the solution will go from dark blue to colourless)
5) Add sodium thiosulfate again, dropwise until the blue colour disappears (blue to colourless)
6) Now calculate the number of moles of Iodine in solution.
107
Give a method that could be used to answer any strucutured or unstrucutured titration calculation
1) Determine the number of moles
2) Using the mole ratio, find out the unknown mole numbers of all your other substances
3) Decide on the amounts that have reacted for known substances, and decide on the amounts of unknowns.
4) Using variations of n = C x V and n = m/Mr, answer the question
108
Define standard electrode potential
The voltage produced when a half-cell is connected to a standard hydrogen half cell under standard conditions
109
What is standard electrode potential measured in?
Volts (V)
110
A more positive standard electrode potential means?
That the chemical will be EASIER to reduce. Therefore, the chemical is reduced and the position of equilibrium lies to the right
111
A more positive standard electrode potential means that the position of equilibrium lies closer to the.....
RIGHT
112
A LESS positive standard electrode potential means?
That the chemical will be HARDER to reduce. Therefore, the chemical is oxidised and the position of equilibrium lies to the left
113
A LESS positive standard electrode potential means that the position of equilibrium lies closer to the.....
LEFT
114
In terms of standard electrode potentials, what are the standard conditions?
Electrode of a pure metal (usually pure platinum) with a wire
298K (25 degrees celsius)
Pressure = 101kPa/1atm
1.00 mol/dm^3 of the chemical in solution
115
In terms of standard electrode potentials, what standard half-cell is used to make comparisons with all other half-cells?
The hydrogen half-cell
H2/H^+ half cell
116
What are the two possible electrode reactions at the hydrogen/half-cell?
2H+ + 2e- ⇌ H2(g) AND H2(g) ⇌ 2H+ + 2e-
117
Why is a pure platinum electrode used?
Inert and it conducts electricity
118
What are the limitations of standard electrode potentials?
Conditions aren't always standard
Rate of reaction may be too slow so it may appear as if it is not even happening
High activation energy. This may even stop the reaction entirely from happening.
119
What is a salt bridge?
A tube that is made out of filter paper soaked in KNO3 or NH4NO3.
120
Why is a salt bridge added to the circuit?
It completes it by allowing for the movement of IONS.
121
Why is a salt bridge needed in a reaction?
It allows for the movement of ions
122
Without a salt bridge can the reaction still proceed?
No.
123
Why is a wire needed in a reaction?
It allows for the transfer of electrons.
124
Give the formula for Eθ cell
Eθ cell = Eθ (positive electrode) - Eθ(negative electrode)
125
How can we predict if X will oxidise Y?
Apply Eθ laws
Most positive Eθ value gets reduced
Least positive Eθ value gets oxidised
126
What other method can we use to represent electrochemical cells
Cell diagrams
127
In a cell diagram, what does a single line represent?
A phase change between the atom.ion
128
In a cell diagram, what does double lines represent?
A salt bridge.
129
In a cell-diagram, what is always on the LHS?
The chemical with the MOST NEGATIVE Eθ
130
In a cell-diagram, what is always on the RHS?
The chemical with the MOST POSITIVE Eθ
131
What does a cell-diagram look like?
X atom I X ion II Y ion I Y atom
Where X is the chemical that is oxidised
Where Y is the chemical that is reduced.
132
If a metal is more reactive, the willingness to lose electrons to form a positive metal ion.
Increases
133
More reactive metals have more ......... Eθ value
NEGATIVE
134
If a non-metal is more reactive, the willingness to gain electrons to form a negative ion
Increases
135
More reactive non-metals have more ......... Eθ value
POSITIVE
136
Magnesium has a more negative Eθ value than Zinc. What does this suggest?
As Mg has a more negative Eθ value that Zinc, it suggests that Mg is more eager to loser electrons to form a positive ion. Therefore, Mg is more reactive than Zinc.
137
Chlorine has a more positive Eθ value Bromine. What does this suggest?
As Cl has a more positive Eθ value than Br, it suggests that Cl is more eager to gain electrons to form a negative ion. Therefore, Cl is more reactive than Br.
138
What is meant by an E value?
An electrode potential value that can be measured under non-standard conditions for a half-cell connected to a hydrogen half-cell.
139
What is meant by storage cell?
A storage cell is an electrochemical cell that can be recharged.
140
With storage cells, a redox reaction takes place. When the reaction is finished what happens?
The cell recharges
141
How can we find he recharging equation of a storage cell
Flip the redo equation for the usage of the storage cell.
142
Give examples of storage cells
Nickel battery
Cadmium battery
Lithium ion and lithium polymer batteries in laptops
143
What is meant by a fuel cell?
An electrochemical cell in which a fuel loses electrons at one electrode (fuel is oxidised) and oxygen gains electrons at the other electrode (oxygen is reduced)
144
What happens in a fuel cell?
An fuel (must contain hydrogen e.g. H2 or methanol) is reacted with oxygen to create a voltage.
The energy produced in the reaction is then converted into electrical energy.
145
If ever, when will fuel cells stop working?
When either oxygen or a fuel is no longer supplied
146
When fuel cells create a voltage, what is the waste product?
Water
147
List benefits of using a fuel cell
- Produce no CO, C, SO2 and NO. (Only waste product is water)
- They are more efficient that combustion engines
- Provide high levels of battery life
148
List drawbacks of using a fuel cell
- Production involves the use of toxic chemicals which need to be disposed of once the cell reaches the end of its life spam
- Cells are highly flammable = May result in fires and explosions
- Need to be replaced more regularly than a combustion engine
- Depending on the fuel used, they can produce little or no CO2.
