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Chemistry > Unit 3.6 - Enthalpy changes for solids and solutions > Flashcards

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

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2
Q

Hess’ law

A

The enthalpy of reaction is independent of the pathway taken by the reaction

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3
Q

Which method do we use for Hess’ law?

A

The “route 1 = route 2” method

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4
Q

Exothermic reactions

A

Energy released to the surroundings during the reaction (hot)

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5
Q

Endothermic
reactions

A

Energy absorbed from the surroundings during the reaction (cold)

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6
Q

Enthalpy change form an energy diagram (exothermic or endothermic reactions)

A

ΔH = Ef - Eb

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7
Q

Is ΔH positive or negative for an exothermic reaction?

A

Negative

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8
Q

Is ΔH positive or negative for an endothermic reaction?

A

Positive

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9
Q

ΔHr

A

Enthalpy change of reaction

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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

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11
Q

Most thermodynamically stable form of carbon

A

Graphite

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12
Q

Why is graphite used for carbon in its standard state?

A

Its the most thermodynamically stable form of carbon

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13
Q

Write the standard enthalpy of formation equation for the formation of ethanol

A

2C (s) + 3H2 (g) + 1/2O2 (g) —> CH3CH2OH (l)

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14
Q

In which direction do we make our arrows point with formation data?

A

Always up

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15
Q

When do we always draw the arrows up for an energy cycle?

A

With formation data

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16
Q

Shortcut method for the standard enthalpy of formation

A

Products - reactants

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17
Q

Standard enthalpy of combustion

A

The enthalpy change when one mole of a substance completely combusts in oxygen under standard conditions

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18
Q

Equation for the standard enthalpy of combustion of ethanol

A

CH3CH2OH (l) + 31/2O2 (g) —> 2CO2 (g) + 3H2O (g)

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19
Q

Reactant of combustion reactions

A

O2

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20
Q

Products of every combustion reaction

A

CO2 + H2O

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21
Q

In what form is H2O formed in combustion reactions and how is this represented?

A

Steam
H2O (g)

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22
Q

In which direction do the arrows point with combustion data?

A

Down

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23
Q

Shortcut method for working out the standard enthalpy of combustion

A

Reactants - products

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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

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25
Q

How do we work out bond enthalpies?

A

Bond enthalpies = reactants - products
(BERP)

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26
Q

Is it possible to measure the actual enthalpy, H, of a system?

A

No

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27
Q

What is it not possible to measure for a system?

A

The actual enthalpy, H

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28
Q

What can we do instead since we can’t measure the actual enthalpy, H, of a system?

A

Can measure enthalpy changes, ΔH

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29
Q

How do we measure enthalpy changes?

A

Can compare the enthalpy of a compound with the enthalpy of the elements it is formed from

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30
Q

Under which conditions can we compare the stability of different substances?

A

Under standard conditions

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31
Q

What is the standard enthalpy change of formation, ΔfH?

A

Comparing the enthalpy of a compound with the enthalpy of the elements it is formed from, under standard conditions

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32
Q

Enthalpy change of formation of anything in their standard state + explanation

A

Zero
Zero energy + already in their standard states and are stable

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33
Q

What is zero when elements are in their standard states

A

Enthalpy change of formation

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34
Q

What makes a compound more stable?

A

More negative enthalpy change of formation

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35
Q

If the enthalpy change of formation is negative, what has happened?

A

Energy has been released (is exothermic)

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36
Q

What does a more negative enthalpy change of formation mean?

A

More stable compound

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37
Q

What type of compounds are the most stable?

A

Most negative enthalpy change of formation (most exothermic)

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38
Q

What has happened if the enthalpy change of formation is positive?

A

Energy has been absorbed (endothermic)

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39
Q

If the enthalpy change of formation is more positive, what does this mean?

A

The compound is less stable

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40
Q

What makes a compound less stable?

A

More positive enthalpy change of formation

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41
Q

What type of compounds are the least stable?

A

More positive enthalpy change of formation
Less exothermic

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42
Q

What do reactions with high enthalpy changes of formation often not do and why?

A

Often do not decompose as the process is too slow

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43
Q

Why do reactions with positive enthalpy changes of formations often not decompose?

A

The process is too slow

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44
Q

What is the equation q = mcΔT used for?

A

To measure the heat energy transferred to solution

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45
Q

Equation for measuring the heat energy transferred to solution

A

q = mcΔT

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46
Q

Define the different units of q = mcΔT

A

q = heat (J)
m = mass of solution
c = specific heat capacity
ΔT = change in temperature

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47
Q

What is m the mass of in q = mcΔT?

A

The solution

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48
Q

How do we find ΔT for q = mcΔT?

A

Between extrapolated lines on a graph

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49
Q

Explain how and why ΔT is found from a graph for q = mcΔT

A

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)

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50
Q

How do we get the enthalpy change from q = mcΔc?

A

ΔH = -q/n

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51
Q

Lattice energy

A

The energy which holds the ions together in ionic crystalline solids

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52
Q

What is lattice energy a measure of and why?

A

The stability of the crystal
The more negative the value, the greater the stability of the lattice

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53
Q

What does a more negative lattice energy value mean?

A

Greater stability of the lattice

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54
Q

Can lattice energy be measured directly by experiment?

A

No

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55
Q

2 types of enthalpy changes for lattice energy

A

Enthalpy change of lattice formation
Enthalpy change of lattice breaking

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56
Q

What type of process is the enthalpy change of lattice formation and why?

A

Exothermic - bond formation

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57
Q

What type of process is any bond formation?

A

Exothermic

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58
Q

What type of process is the enthalpy change of lattice breaking and why?

A

Endothermic
Bond breaking

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59
Q

What type of reaction is any bond breaking?

A

Endothermic

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60
Q

Enthalpy change of lattice breaking

A

The energy required to change 1 mole of a crystalline solid into gaseous ions under standard conditions

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61
Q

Enthalpy change of formation

A

The enthalpy change when one mole of a substance forms from its constituent elements in their standard states under standard conditions

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62
Q

Equation for the enthalpy change of formation of NaCl

A

Na (s) + 1/2Cl2 (g) —> NaCl (s)

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63
Q

Enthalpy change of atomisation

A

The enthalpy change which occurs when a substance in its standard state under standard conditions is changes into 1 mol of gaseous atoms

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64
Q

What happens for metals during enthalpy change of atomisation?

A

A solid is changes into a gas

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65
Q

Solid changed into a gas

A

Sublimation

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66
Q

What does sublimation involve and what does this make it?

A

Bond breaking
Endothermic

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67
Q

Equation for Na undergoing sublimation (enthalpy change of atomisation)

A

Na (s) —> Na (g)

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68
Q

What basically happens during enthalpy change of atomisation

A

Standard state —> gaseous atoms

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69
Q

Enthalpy change of solution

A

The Enthalpy change when 1 mole of a substance is dissolved in water

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70
Q

Equation for enthalpy change of solution of NaCl

A

NaCl (s) + water —> Na+ (aq) + Cl- (aq)

71
Q

Enthalpy change of hydration

A

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
Q

Equation for the enthalpy change of hydration of Na+

A

Na+ (g) + water —> Na+ (aq)

73
Q

1st ionisation energy

A

Energy change when 1 mole of gaseous atoms are changed into 1 mole of cations, each atom losing one electron

74
Q

Equation for the 1st ionisation energy of Na

A

Na (g) —> Na+ (g)

75
Q

What type of process is 1st ionisation energy

A

Endothermic

76
Q

What are required for group 2 elements in terms of ionisation energy? Explain

A

Successive ionisation energies
I.e - for group 2, we need second and first ionisation energy

77
Q

Bond energy

A

Enthalpy change when 1 mole of gaseous molecules and converted into gaseous atoms

78
Q

What type of process is bond energy?

A

Endothermic

79
Q

Equation for the bond energy of Cl2

A

Cl2 (g) —> 2Cl (g)

80
Q

Symbol for bond energy enthalpy change

A

ΔHD

81
Q

In the below equation, what enthalpy change needs to be used and why?
Cl2 (g) —> 2Cl (g)

A

As only 1/2 mole of chlorine molecules are converted into atoms, 1/2 of the enthalpy change is required

82
Q

Electron affinity

A

The energy required to change 1 mole of gaseous atoms into gaseous anions, each atom gaining 1 electron

83
Q

What type of process is electrons affinity?

A

Exothermic

84
Q

what can Cl be classified as?

A

A radical

85
Q

Equation for the electron affinity of Cl

A

Cl (g) + e- —> Cl-(g)

86
Q

When an anion such as O^2- is formed, what is the overall process? (Electron affinity)

A

Endothermic

87
Q

Why is the overall process endothermic when an anion such as O^2- is formed? (Electron affinity)

A

Since the second electron is added to a negative species

88
Q

Equations for the 1st and 2nd electron affinities of oxygen

A

1st electron affinity of oxygen:
O (g) + e- —> O- (g)

2nd electron affinity of oxygen:
O- (g) + e- —> O2- (g)

89
Q

What type of reaction is 1st electron affinity always?

A

Exothermic

90
Q

What type of reaction is 2nd electron affinity always?

A

Endothermic

91
Q

What are the 2 steps involved when an ionic solid dissolves in water?

A

1.) lattice breaking
2.) hydration

92
Q

Lattice breaking

A

The separation of the ions

93
Q

Hydration

A

Combining the ions with water molecules

94
Q

What type of reaction is any lattice breaking?

A

Endothermic

95
Q

What type of reaction is any hydration reaction?

A

Exothermic

96
Q

Enthalpy change of solution

A

The enthalpy change when 1 mole of a substance is dissolved in water

97
Q

Equation for the enthalpy change of solution of NaCl

A

NaCl (s) + water —> Na+ (aq) + Cl- (aq)

98
Q

What type of reactions are enthalpy changes of solution

A

Physical processes, not chemical reactions

99
Q

What can be drawn to show how the enthalpy change of solution is related to lattice energies and hydration energies?

A

An energy cycle

100
Q

What does an energy cycle of the enthalpy change of solution also include?

A

Lattice energies
Hydration energies

101
Q

What can we apply to energy cycles?

A

Hess’ law

102
Q

Enthalpy change of solution equation

A

ΔH solution = ΔH lattice + ΔH hydration

103
Q

Why would an enthalpy change of solution be positive?

A

Due to entropy change

104
Q

When will a substance dissolve in water?

A

If the process is energetically favourable

105
Q

Which conditions need to be true for an ionic solid to dissolve in water?

A

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
Q

Why is he enthalpy change of solution for NaCl small compared to the values of lattice energy and hydration energy?

A

Since the values of lattice energy and hydration energy are large but are fairly similar in value

107
Q

What is the actual value of hydration energy dependent on?

A

1.) hydration energy increases as the charge on the ion increases
2.) hydration energy decreases as the ionic radius increases

108
Q

Give an example of how hydration energy increases as the charge on the ion increases

A

Na+ > Mg2+ > Al3+

109
Q

Why does hydration energy increase as the charge on the ion increases?

A

More attraction to water molecules

110
Q

Why does hydration energy decrease as the ionic radius increases?

A

Less attraction between the nucleus and the negative end of the water molecule = shielding

111
Q

What needs to be true for an ionic solid to dissolve in terms of energy?

A

The hydration energy needs to be similar to or larger than the lattice energy

112
Q

Under which situation will an ionic substance not dissolve in water?

A

If the lattice energy is somewhat larger than the hydration energy

113
Q

What type of value should the enthalpy change of solution be?

A

Negative (exothermic)

114
Q

When will we know if a substance is soluble and will dissolve?

A

When it’s enthalpy change of solution is exothermic

115
Q

What does an exothermic enthalpy change of solution show us?

A

That the substance is soluble and will dissolve

116
Q

If the enthalpy change of solution of a substance is slightly endothermic, will it still dissolve?

A

Yes

117
Q

Equation for lattice breaking of CaCl

A

CaCl (s) —> 2Cl- (g) + Ca2+ (g)

118
Q

Equation for hydration of CaCl

A

(Lattice breaking would have occurred first to form Ca2+ (g))
Ca2+ (g) + 2Cl- + q —> CaCl2 (g)

119
Q

What is the opposite of 1st ionisation energy? Explain

A

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
Q

What is the general rule for what makes a compound more stable?

A

If its standard enthalpy change of formation is exothermic (i.e - negative)

121
Q

What does a more exothermic/negative enthalpy change of formation of a compound mean?

A

That it’s more stable

122
Q

What does an endothermic enthalpy change of formation imply?

A

That the compound is likely to be unstable

123
Q

What type of compounds are most likely to be unstable?

A

Those with endothermic enthalpy changes of formation

124
Q

Is it always true that an exothermic enthalpy change of formation means a more stable compound?

A

No - it’s a qualitative effect only so only use it generally

125
Q

What has to be done to a metal in order to extract it from its ore?

A

Has to be reduced

126
Q

How are metals normally found in nature?

A

Combines with other elements

127
Q

Examples of metals found combined with other elements in nature

A

Zinc sulphide, iron (III) oxide, magnesium chloride

128
Q

Example of an uncombined metal

A

Gold

129
Q

What type of a metal is gold?

A

Uncombined

130
Q

What does the method used to extract a metal from its ore largely depend on?

A

The amount of energy required to reduce the metal ion to the metal

131
Q

How can we get an indication as to the ease or difficulty of the extraction of a metal from its ore?

A

From the value of the enthalpy change of formation of its oxide

132
Q

Explain how we would identify metals which are easily extracted from their oxides

A

Endothermic or small exothermic values for the enthalpy change of formation of the oxide

133
Q

What does an endothermic or small exothermic value for the enthalpy change of formation of an oxide mean in terms of its extraction?

A

Metal easily extracted from its oxide

134
Q

2 metals which are found naturally occurring as the metal

A

Gold
Silver

135
Q

Explain how we would identify metals which should be extracted from their compounds by carbon reduction

A

Large exothermic values for the enthalpy change of formation of the oxide

136
Q

What does a large exothermic value for the enthalpy change of formation of the oxide mean in terms of the extraction of a metal?

A

Metals which are extracted from their compounds by carbon reduction

137
Q

Which metals are extracted from their compounds by carbon reduction?

A

Metals above copper in the electrochemical series

138
Q

Equation for the extraction of Pb from PbO by carbon reduction

A

PbO (s) + C (s) —> Pb (s) + CO (g)

139
Q

What do metals with very large exothermic values for the enthalpy change of formation need to be extracted from their compounds?

A

Considerable amounts of energy

140
Q

What’s the normal method used to extract metals with very large exothermic values for the enthalpy change of formation of the oxide?

A

Electrolysis

141
Q

When is electrolysis used to extract a metal?

A

Metals with very large exothermic values for the enthalpy change of formation of the oxide

142
Q

Examples of metals extracted by electrolysis

A

Aluminium, calcium, sodium

143
Q

What makes a compound stable?

A

Strong bonds

144
Q

What must be done to break strong bonds?

A

Lots of energy must be absorbed (endothermic)

145
Q

Why does a more negative enthalpy change of formation mean that the compound is more stable? Explain

A

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
Q

What are Born-Haber cycles?

A

Energy cycles with several steps that are used to calculate the energy of formation of ionic compounds from elements in their standard states

147
Q

What are Born-Haber cycles used for?

A

To calculate the energy of formation of ionic compounds from elements in their standard states

148
Q

What steps do Born-Haber cycles include?

A

Many of the energy transfers

149
Q

What can we apply to Born Haber cycles?

A

Hess’ law

150
Q

What do upward and downward facing arrows represent on Born-Haber cycles?

A

Upwards = endothermic
Downwards = exothermic

151
Q

What are the 2 routes shown on a Born-Haber cycle?

A

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
Q

Order of the steps in the longer process of a Born-Haber cycle

A

Enthalpy of atomisation
Ionisation energy
Enthalpy of atomisation/bond dissociation
Electron affinity
Lattice formation

153
Q

What does a more exothermic enthalpy change of lattice formation lead to?

A

A more stable ionic compound

154
Q

How do you know if an ionic compound is more stable from lattice formation values?

A

More exothermic = more stable ionic compound

155
Q

Enthalpy occurs atomisation of Li equatiion

A

Li (s) —> Li (g)

156
Q

Equation for the first ionisation energy of lithium

A

Li (g) —> Li+ (g) + e-

157
Q

Equation for the enthalpy of atomisation/bond dissociation of fluorine

A

1/2F2 (g) —> F (g)

158
Q

Equation for electron affinity of fluorine

A

F (g) —> F- (g)

159
Q

Equation for the lattice formation of LiF

A

Li+ (g) + F- (g) —> LiF (s)

160
Q

Important concept for Hess’ law

A

Route 1 = route 2

161
Q

What is the enthalpy change of formation equal to on most Born-Haber cycles?

A

Atomisation + ionisation + bond dissociation/atomisation + electron affinity + lattice formation

162
Q

What could we be given when referring to ionisation?

A

Total, 1st, 2nd or 3rd

163
Q

What do we do if we’re given the electron affinity for 1 Cl- ion and need it for Cl2?

A

Multiply the enthalpy change by 2

164
Q

What do we do if we’re given the atomisation/bond dissociation for Cl2, but we only need it for Cl-?

A

Divide the enthalpy change by 2

165
Q

How are 1st and 2nd electron affinities different and how could this affect a Born-Haber cycle?

A

1st electron affinity = exothermic
2nd electron affinity = endothermic
If it were endothermic, it would have the arrow facing in the opposite direction

166
Q

Explain why the second electron affinity of oxygen is positive

A

The second electron is added to a negative species so its endothermic since energy is needed to do this

167
Q

How do we apply Hess’ law to a Born-Haber cycle with an arrow facing in the opposite direction for the 2nd electron affinity?

A

It’s still in the same direction as everything else except for the enthalpy of formation

168
Q

What do we need to not get confused between with Born-Haber cycles?

A

ΔHf between…
Enthalpy of formation (ΔHf)
Lattice formation (ΔHL)

169
Q

Which values might we need to multiply or divide by 2 when working with Born-Haber cycles?

A

Electron affinity
Ionisation
Atomisation/bond dissociation
Bond energy

170
Q

What do you do if you’re given a lattice breaking value with Born-Haber cycles? Why?

A

Change the + to -
We want lattice formation values

171
Q

What do we need to remember to do when writing about enthalpy changes of atomisation?

A

Write it for one atom of that element
E.g - Cl2 = 242
So, write enthalpy change of atomisation as 121kJmol-1

172
Q

What happens to ionisation energies with successive ionisation energies?

A

Increase for successive ionisation energies

173
Q

What will have the highest ionisation energy value? 2nd or 1st?

A

2nd

174
Q

Describe how the enthalpy changes need to be for an ionic solid to dissolve

A

Enthalpy change of hydration needs to be more exothermic than the enthalpy change of lattice making

Chemistry (39 decks)

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