Unit 3.6 - Enthalpy changes for solids and solutions Flashcards
1
Q
Principle of conservation of energy
A
Energy cannot be created or destroyed, only transformed from one form to another
2
Q
Hess’ law
A
The enthalpy of reaction is independent of the pathway taken by the reaction
3
Q
Which method do we use for Hess’ law?
A
The “route 1 = route 2” method
4
Q
Exothermic reactions
A
Energy released to the surroundings during the reaction (hot)
5
Q
Endothermic
reactions
A
Energy absorbed from the surroundings during the reaction (cold)
6
Q
Enthalpy change form an energy diagram (exothermic or endothermic reactions)
A
ΔH = Ef - Eb
7
Q
Is ΔH positive or negative for an exothermic reaction?
A
Negative
8
Q
Is ΔH positive or negative for an endothermic reaction?
A
Positive
9
Q
ΔHr
A
Enthalpy change of reaction
10
Q
Standard enthalpy of formation
A
The enthalpy change when one mole of a substance is formed from its constituent elements in their standard states under standard conditions
11
Q
Most thermodynamically stable form of carbon
A
Graphite
12
Q
Why is graphite used for carbon in its standard state?
A
Its the most thermodynamically stable form of carbon
13
Q
Write the standard enthalpy of formation equation for the formation of ethanol
A
2C (s) + 3H2 (g) + 1/2O2 (g) —> CH3CH2OH (l)
14
Q
In which direction do we make our arrows point with formation data?
A
Always up
15
Q
When do we always draw the arrows up for an energy cycle?
A
With formation data
16
Q
Shortcut method for the standard enthalpy of formation
A
Products - reactants
17
Q
Standard enthalpy of combustion
A
The enthalpy change when one mole of a substance completely combusts in oxygen under standard conditions
18
Q
Equation for the standard enthalpy of combustion of ethanol
A
CH3CH2OH (l) + 31/2O2 (g) —> 2CO2 (g) + 3H2O (g)
19
Q
Reactant of combustion reactions
A
O2
20
Q
Products of every combustion reaction
A
CO2 + H2O
21
Q
In what form is H2O formed in combustion reactions and how is this represented?
A
Steam
H2O (g)
22
Q
In which direction do the arrows point with combustion data?
A
Down
23
Q
Shortcut method for working out the standard enthalpy of combustion
A
Reactants - products
24
Q
Average bond enthalpy
A
The amount of energy required to break one mole of bonds of a particular type between two atoms in gaseous state
25
How do we work out bond enthalpies?
Bond enthalpies = reactants - products
(BERP)
26
Is it possible to measure the actual enthalpy, H, of a system?
No
27
What is it not possible to measure for a system?
The actual enthalpy, H
28
What can we do instead since we can’t measure the actual enthalpy, H, of a system?
Can measure enthalpy changes, ΔH
29
How do we measure enthalpy changes?
Can compare the enthalpy of a compound with the enthalpy of the elements it is formed from
30
Under which conditions can we compare the stability of different substances?
Under standard conditions
31
What is the standard enthalpy change of formation, ΔfH?
Comparing the enthalpy of a compound with the enthalpy of the elements it is formed from, under standard conditions
32
Enthalpy change of formation of anything in their standard state + explanation
Zero
Zero energy + already in their standard states and are stable
33
What is zero when elements are in their standard states
Enthalpy change of *formation*
34
What makes a compound more stable?
More negative enthalpy change of formation
35
If the enthalpy change of formation is negative, what has happened?
Energy has been released (is exothermic)
36
What does a more negative enthalpy change of formation mean?
More stable compound
37
What type of compounds are the most stable?
Most negative enthalpy change of formation (most exothermic)
38
What has happened if the enthalpy change of formation is positive?
Energy has been absorbed (endothermic)
39
If the enthalpy change of formation is more positive, what does this mean?
The compound is less stable
40
What makes a compound less stable?
More positive enthalpy change of formation
41
What type of compounds are the least stable?
More positive enthalpy change of formation
Less exothermic
42
What do reactions with high enthalpy changes of formation often not do and why?
Often do not decompose as the process is too slow
43
Why do reactions with positive enthalpy changes of formations often not decompose?
The process is too slow
44
What is the equation q = mcΔT used for?
To measure the heat energy transferred to solution
45
Equation for measuring the heat energy transferred to solution
q = mcΔT
46
Define the different units of q = mcΔT
q = heat (J)
m = mass of *solution*
c = specific heat capacity
ΔT = change in temperature
47
What is m the mass of in q = mcΔT?
The *solution*
48
How do we find ΔT for q = mcΔT?
Between extrapolated lines on a graph
49
Explain how and why ΔT is found from a graph for q = mcΔT
ΔT is between extrapolated lines on a graph
When a certain solid is added to solution, it causes a temperature rise.
We don’t record the temperature immediately after adding it, but after stirring for some time we get a graph similar to the one seen in the notes (imagine some low plots in a line, then a big jump to some high plots - the difference between these is ΔT)
50
How do we get the enthalpy change from q = mcΔc?
ΔH = -q/n
51
Lattice energy
The energy which holds the ions together in ionic crystalline solids
52
What is lattice energy a measure of and why?
The stability of the crystal
The more negative the value, the greater the stability of the lattice
53
What does a more negative lattice energy value mean?
Greater stability of the lattice
54
Can lattice energy be measured directly by experiment?
No
55
2 types of enthalpy changes for lattice energy
Enthalpy change of lattice formation
Enthalpy change of lattice breaking
56
What type of process is the enthalpy change of lattice formation and why?
Exothermic - bond formation
57
What type of process is any bond formation?
Exothermic
58
What type of process is the enthalpy change of lattice breaking and why?
Endothermic
Bond breaking
59
What type of reaction is any bond breaking?
Endothermic
60
Enthalpy change of lattice breaking
The energy required to change 1 mole of a crystalline solid into gaseous ions under standard conditions
61
Enthalpy change of formation
The enthalpy change when one mole of a substance forms from its constituent elements in their standard states under standard conditions
62
Equation for the enthalpy change of formation of NaCl
Na (s) + 1/2Cl2 (g) —> NaCl (s)
63
Enthalpy change of atomisation
The enthalpy change which occurs when a substance in its standard state under standard conditions is changes into 1 mol of gaseous atoms
64
What happens for metals during enthalpy change of atomisation?
A solid is changes into a gas
65
Solid changed into a gas
Sublimation
66
What does sublimation involve and what does this make it?
Bond breaking
Endothermic
67
Equation for Na undergoing sublimation (enthalpy change of atomisation)
Na (s) —> Na (g)
68
What basically happens during enthalpy change of atomisation
Standard state —> gaseous atoms
69
Enthalpy change of solution
The Enthalpy change when 1 mole of a substance is dissolved in water
70
Equation for enthalpy change of solution of NaCl
NaCl (s) + water —> Na+ (aq) + Cl- (aq)
71
Enthalpy change of hydration
The energy released when 1 mole of ions in the gaseous state are converted into hydrated ions (1 mole of an ionic compound in solution)
72
Equation for the enthalpy change of hydration of Na+
Na+ (g) + water —> Na+ (aq)
73
1st ionisation energy
Energy change when 1 mole of gaseous atoms are changed into 1 mole of cations, each atom losing one electron
74
Equation for the 1st ionisation energy of Na
Na (g) —> Na+ (g)
75
What type of process is 1st ionisation energy
Endothermic
76
What are required for group 2 elements in terms of ionisation energy? Explain
Successive ionisation energies
I.e - for group 2, we need second and first ionisation energy
77
Bond energy
Enthalpy change when 1 mole of gaseous molecules and converted into gaseous atoms
78
What type of process is bond energy?
Endothermic
79
Equation for the bond energy of Cl2
Cl2 (g) —> 2Cl (g)
80
Symbol for bond energy enthalpy change
ΔHD
81
In the below equation, what enthalpy change needs to be used and why?
Cl2 (g) —> 2Cl (g)
As only 1/2 mole of chlorine molecules are converted into atoms, 1/2 of the enthalpy change is required
82
Electron affinity
The energy required to change 1 mole of gaseous atoms into gaseous anions, each atom gaining 1 electron
83
What type of process is electrons affinity?
Exothermic
84
what can Cl be classified as?
A radical
85
Equation for the electron affinity of Cl
Cl (g) + e- —> Cl-(g)
86
When an anion such as O^2- is formed, what is the overall process? (Electron affinity)
Endothermic
87
Why is the overall process endothermic when an anion such as O^2- is formed? (Electron affinity)
Since the second electron is added to a negative species
88
Equations for the 1st and 2nd electron affinities of oxygen
1st electron affinity of oxygen:
O (g) + e- —> O- (g)
2nd electron affinity of oxygen:
O- (g) + e- —> O2- (g)
89
What type of reaction is 1st electron affinity always?
Exothermic
90
What type of reaction is 2nd electron affinity always?
Endothermic
91
What are the 2 steps involved when an ionic solid dissolves in water?
1.) lattice breaking
2.) hydration
92
Lattice breaking
The separation of the ions
93
Hydration
Combining the ions with water molecules
94
What type of reaction is any lattice breaking?
Endothermic
95
What type of reaction is any hydration reaction?
Exothermic
96
Enthalpy change of solution
The enthalpy change when 1 mole of a substance is dissolved in water
97
Equation for the enthalpy change of solution of NaCl
NaCl (s) + water —> Na+ (aq) + Cl- (aq)
98
What type of reactions are enthalpy changes of solution
Physical processes, not chemical reactions
99
What can be drawn to show how the enthalpy change of solution is related to lattice energies and hydration energies?
An energy cycle
100
What does an energy cycle of the enthalpy change of solution also include?
Lattice energies
Hydration energies
101
What can we apply to energy cycles?
Hess’ law
102
Enthalpy change of solution equation
ΔH solution = ΔH lattice + ΔH hydration
103
Why would an enthalpy change of solution be positive?
Due to entropy change
104
When will a substance dissolve in water?
If the process is energetically favourable
105
Which conditions need to be true for an ionic solid to dissolve in water?
The strength of the ion-dipole forces between the ions and water molecules must be similar to or stronger than the ionic forces in the lattice
106
Why is he enthalpy change of solution for NaCl small compared to the values of lattice energy and hydration energy?
Since the values of lattice energy and hydration energy are large but are fairly similar in value
107
What is the actual value of hydration energy dependent on?
1.) hydration energy increases as the charge on the ion increases
2.) hydration energy decreases as the ionic radius increases
108
Give an example of how hydration energy increases as the charge on the ion increases
Na+ > Mg2+ > Al3+
109
Why does hydration energy increase as the charge on the ion increases?
More attraction to water molecules
110
Why does hydration energy decrease as the ionic radius increases?
Less attraction between the nucleus and the negative end of the water molecule = shielding
111
What needs to be true for an ionic solid to dissolve in terms of energy?
The hydration energy needs to be similar to or larger than the lattice energy
112
Under which situation will an ionic substance not dissolve in water?
If the lattice energy is somewhat larger than the hydration energy
113
What type of value should the enthalpy change of solution be?
Negative (exothermic)
114
When will we know if a substance is soluble and will dissolve?
When it’s enthalpy change of solution is exothermic
115
What does an exothermic enthalpy change of solution show us?
That the substance is soluble and will dissolve
116
If the enthalpy change of solution of a substance is slightly endothermic, will it still dissolve?
Yes
117
Equation for lattice breaking of CaCl
CaCl (s) —> 2Cl- (g) + Ca2+ (g)
118
Equation for hydration of CaCl
(Lattice breaking would have occurred first to form Ca2+ (g))
Ca2+ (g) + 2Cl- + q —> CaCl2 (g)
119
What is the opposite of 1st ionisation energy? Explain
Electron affinity
1st ionisation energy = 1mol of gaseous atoms changes into cations, with each atom losing 1 electron
Electron affinity = 1mol of gaseous atoms changes into anions, with each atom gaining 1 electron
120
What is the general rule for what makes a compound more stable?
If its standard enthalpy change of formation is exothermic (i.e - negative)
121
What does a more exothermic/negative enthalpy change of formation of a compound mean?
That it’s more stable
122
What does an endothermic enthalpy change of formation imply?
That the compound is likely to be unstable
123
What type of compounds are most likely to be unstable?
Those with endothermic enthalpy changes of formation
124
Is it always true that an exothermic enthalpy change of formation means a more stable compound?
No - it’s a qualitative effect only so only use it generally
125
What has to be done to a metal in order to extract it from its ore?
Has to be reduced
126
How are metals normally found in nature?
Combines with other elements
127
Examples of metals found combined with other elements in nature
Zinc sulphide, iron (III) oxide, magnesium chloride
128
Example of an uncombined metal
Gold
129
What type of a metal is gold?
Uncombined
130
What does the method used to extract a metal from its ore largely depend on?
The amount of energy required to reduce the metal ion to the metal
131
How can we get an indication as to the ease or difficulty of the extraction of a metal from its ore?
From the value of the enthalpy change of formation of its oxide
132
Explain how we would identify metals which are easily extracted from their oxides
Endothermic or small exothermic values for the enthalpy change of formation of the oxide
133
What does an endothermic or small exothermic value for the enthalpy change of formation of an oxide mean in terms of its extraction?
Metal easily extracted from its oxide
134
2 metals which are found naturally occurring as the metal
Gold
Silver
135
Explain how we would identify metals which should be extracted from their compounds by carbon reduction
Large exothermic values for the enthalpy change of formation of the oxide
136
What does a large exothermic value for the enthalpy change of formation of the oxide mean in terms of the extraction of a metal?
Metals which are extracted from their compounds by carbon reduction
137
Which metals are extracted from their compounds by carbon reduction?
Metals above copper in the electrochemical series
138
Equation for the extraction of Pb from PbO by carbon reduction
PbO (s) + C (s) —> Pb (s) + CO (g)
139
What do metals with *very* large exothermic values for the enthalpy change of formation need to be extracted from their compounds?
Considerable amounts of energy
140
What’s the normal method used to extract metals with very large exothermic values for the enthalpy change of formation of the oxide?
Electrolysis
141
When is electrolysis used to extract a metal?
Metals with very large exothermic values for the enthalpy change of formation of the oxide
142
Examples of metals extracted by electrolysis
Aluminium, calcium, sodium
143
What makes a compound stable?
Strong bonds
144
What must be done to break strong bonds?
Lots of energy must be absorbed (endothermic)
145
Why does a more negative enthalpy change of formation mean that the compound is more stable? Explain
Stable compound = strong bonds
Strong bonds = lots of energy must be absorbed to break the bonds (endothermic)
Conversely, in forming the compound, the same amount of energy is given out
So, the more negative the enthalpy change of formation of the compound, the more stable it will be
146
What are Born-Haber cycles?
Energy cycles with several steps that are used to calculate the energy of formation of ionic compounds from elements in their standard states
147
What are Born-Haber cycles used for?
To calculate the energy of formation of ionic compounds from elements in their standard states
148
What steps do Born-Haber cycles include?
Many of the energy transfers
149
What can we apply to Born Haber cycles?
Hess’ law
150
What do upward and downward facing arrows represent on Born-Haber cycles?
Upwards = endothermic
Downwards = exothermic
151
What are the 2 routes shown on a Born-Haber cycle?
Route 1 = the energy of formation from the elements in their standard states
Route 2 = the series of steps leading to the formation of the same thing
152
Order of the steps in the longer process of a Born-Haber cycle
Enthalpy of atomisation
Ionisation energy
Enthalpy of atomisation/bond dissociation
Electron affinity
Lattice formation
153
What does a more exothermic enthalpy change of lattice formation lead to?
A more stable ionic compound
154
How do you know if an ionic compound is more stable from lattice formation values?
More exothermic = more stable ionic compound
155
Enthalpy occurs atomisation of Li equatiion
Li (s) —> Li (g)
156
Equation for the first ionisation energy of lithium
Li (g) —> Li+ (g) + e-
157
Equation for the enthalpy of atomisation/bond dissociation of fluorine
1/2F2 (g) —> F (g)
158
Equation for electron affinity of fluorine
F (g) —> F- (g)
159
Equation for the lattice formation of LiF
Li+ (g) + F- (g) —> LiF (s)
160
Important concept for Hess’ law
Route 1 = route 2
161
What is the enthalpy change of formation equal to on most Born-Haber cycles?
Atomisation + ionisation + bond dissociation/atomisation + electron affinity + lattice formation
162
What could we be given when referring to ionisation?
Total, 1st, 2nd or 3rd
163
What do we do if we’re given the electron affinity for 1 Cl- ion and need it for Cl2?
Multiply the enthalpy change by 2
164
What do we do if we’re given the atomisation/bond dissociation for Cl2, but we only need it for Cl-?
Divide the enthalpy change by 2
165
How are 1st and 2nd electron affinities different and how could this affect a Born-Haber cycle?
1st electron affinity = exothermic
2nd electron affinity = endothermic
If it were endothermic, it would have the arrow facing in the opposite direction
166
Explain why the second electron affinity of oxygen is positive
The second electron is added to a negative species so its endothermic since energy is needed to do this
167
How do we apply Hess’ law to a Born-Haber cycle with an arrow facing in the opposite direction for the 2nd electron affinity?
It’s still in the same direction as everything else except for the enthalpy of formation
168
What do we need to not get confused between with Born-Haber cycles?
ΔHf between…
Enthalpy of formation (ΔHf)
Lattice formation (ΔHL)
169
Which values might we need to multiply or divide by 2 when working with Born-Haber cycles?
Electron affinity
Ionisation
Atomisation/bond dissociation
Bond energy
170
What do you do if you’re given a lattice breaking value with Born-Haber cycles? Why?
Change the + to -
We want lattice formation values
171
What do we need to remember to do when writing about enthalpy changes of atomisation?
Write it for one atom of that element
E.g - Cl2 = 242
So, write enthalpy change of atomisation as 121kJmol-1
172
What happens to ionisation energies with successive ionisation energies?
Increase for successive ionisation energies
173
What will have the highest ionisation energy value? 2nd or 1st?
2nd
174
Describe how the enthalpy changes need to be for an ionic solid to dissolve
Enthalpy change of hydration needs to be more exothermic than the enthalpy change of lattice making
