Energetics I Flashcards
1
Q
What is enthalpy, ΔE?
A
A measure of energy content of the substance
2
Q
What does lower enthalpy indicate?
A
Greater energetic stability
3
Q
What is enthalpy change, ΔE
A
Change in energy content at constant pressure
4
Q
Is ΔH positive or negative for exothermic reaction and do products or reactants have higher energy?
A
- products have a lower energy than reactants
- ΔH = H(products) - H(reactants), ΔH < 0
5
Q
What happens to the temperature of the surroundings in an exothermic reaction?
A
- heat is released to the surroundings
- temperature of surroundings rises
6
Q
Is ΔH positive or negative for endothermic reaction and do products or reactants have higher energy?
A
- products have higher energy than reactants
- ΔH = H(products) - H(reactants), ΔH > 0
7
Q
What happens to the temperature of the surroundings in an endothermic reaction?
A
- heat is absorbed from the surroundings
- temperature of the surrounding drop
8
Q
What is the system and what is the surroundings for calorimeter?
A
System: reaction
Surroundings: aqueous medium and calorimeter
9
Q
What happens in calorimeter in exothermic reaction?
A
- heat evolved by reaction = heat gained by calorimeter + heat gained by aqueous medium
- thus temperature of reaction mixture increases
10
Q
What happens in calorimeter in endothermic reaction?
A
- heat absorbed by the reaction = heat lost by calorimeter + heat loss by aqueous medium
- temperature of reaction mixture decreases
11
Q
What is the formula for heat evolved/absorbed?
A
Q = mcΔT
12
Q
For dilute solutions, what can the density and the specific heat capacity of the solution be assumed as?
A
- density = density of pure water = 1.00g cm⁻³
- specific heat capacity = 4.18 J g⁻¹ K⁻¹
13
Q
What are the assumptions made in calorimetry experiments?
A
- heat capacity of the calorimeter vessel is negligible, thus there is no heat gained/loss to the calorimeter
- no heat exchange with the surroundings
14
Q
What is specific heat capacity, c?
A
amount of heat required to raise the temperature of 1g of the substance by 1K
15
Q
What is heat capacity, C?
A
heat required to raise the temperature of a given quantity of the substance by 1K
16
Q
What is standard conditions°?
A
- temperature: 298K or 25 degree celsius
- 1 bar or 1.00 × 10⁵ Pa
- most stable physical state (s, l or g) at 298K and 1 bar
- 1 mol dm⁻³
17
Q
What is standard enthalpy change of reaction, ΔHᵣ°?
A
Standard enthalpy change of reaction is the enthalpy change when molar quantities of reactants, as specified by the balanced chemical equation, react to form products under standard conditions
18
Q
What does ‘per mole’ in the unit of ΔH° imply?
A
per mole of equation
19
Q
What does ΔHᵣ° depend on?
A
- magnitude of ΔHᵣ is directly proportional to number of reactants consumed (eg. if 2a + 2b → c, ΔHᵣ×2)
- enthalpy change is equal in magnitude but opposite in sign for the reverse reaction
- enthalpy change for reaction depends on physical state of reactants and products
20
Q
What is a thermochemical equation?
A
A balanced stoichiometric equation, including state symbols for the reaction with ΔH values
21
Q
What is standard enthalphy change of atomisation, ΔHₐₜ°?
A
The standard enthalpy change of atomisation of an element is the enthalpy change when one mole of gaseous atoms is formed from the element in its standard state under standard conditions
22
Q
What is the sign of ΔHₐₜ° values?
A
Always positive because atomisation involves bond breaking which takes in energy
23
Q
What does greater ΔHₐₜ° indicate?
A
The higher the ΔHₐₜ° value, the greater the strength of the bond broken
24
Q
What does ΔHₐₜ include for liquids?
A
enthalpy change of vapourisation, ΔHᵥₐₚ
25
What does ΔHₐₜ include for solids?
enthalpy change of fusion, ΔH(fus) + ΔHᵥₐₚ
26
What is the standard enthalpy change of formation, ΔHf°?
Enthalpy change when **one mole of the substance** is formed from its **constituent elements** in their standard states under standard conditions
27
What is ΔHf° of an element in its standard state?
zero
28
What is ΔHf° a measure of?
Measure of its energetic stability relative to its constituent elements
29
What does ΔHf° < 0 mean?
* Reaction is more exothermic
* substance is more stable relative to its constituent elements
30
What is standard enthalpy change of combustion, ΔHc°?
Enthalpy change when **one mole of the substance** is completely burnt in oxygen under standard conditions
31
What is the sign of ΔHc°?
Always negative because combustion is an exothermic reaction
32
What can ΔHc° values be used for?
* It can be used to compare the energy values of fuels and foods
* the more exothermic the combustion reaction, the higher the energy value of the fuel
33
What are the improvements/ modifications you can suggest for calorimetry experiments?
* minimise heat loss by using thermal lagging on the calorimeter
* use wind shield to reduce draught
* use a lid to prevent heat loss by convection
34
What is first ionisation energy?
First ionisation energy of an element is the enthalpy change when one mole of electrons is removed from **one mole of gaseous atoms**
35
What is second ionisation energy?
Second ionisation energy of an element is the enthalpy change when one mole of electrons is removed from **one mole of gaseous, singly charged cations of the element**
36
What is the sign for ionisation energy?
- Always positive because energy is required to overcome the forces of attraction between the nucleus and the electron to be removed
- reaction always endothermic
37
What is first electron affinity?
Enthalpy change when **one mole of gaseous atoms accepts one mole of electrons**
38
What is second electron affinity?
Enthalpy change when **one mole of gaseous, singly charged anions** accepts one mole of electrons
39
What is electron affinity a measure of?
* Attraction between the incoming electron and the nucleus
* the stronger the attraction, the more energy is released
40
What is the sign of the first electron affinity of electronegative elements?
- negative
- the more electronegative the element, the more negative the value of first electron affinity
41
What is the sign of second and successive electron affinities?
* Usually positive as energy is required to overcome the repulsion between anion and the electron to be added, both of which is negatively charged
* endothermic
42
What is lattice energy?
Enthalpy change when **one mole of solid ionic compound** is formed from its **constituent gaseous ions** under standard conditions
43
What is the sign of lattice energy?
- Always negative
- process is exothermic as energy is evolved when oppositely charged ions come together to form ionic bonds
44
What is lattice energy a measure of?
- strength of ionic bonds in an ionic compound
- the more negative the enthalpy change, the more exothermic the reaction, stronger the ionic bond
45
What is the formula for lattice energy?
LE ∝ |Z⁺ × Z⁻ / r⁺ + r⁻|
Z: charges on ions
r: ionic radius
46
What is bond energy?
**Average** enthalpy change when **one mole of covalent bonds between atoms in gaseous molecules is broken**
47
What is bond dissociation energy?
Enthalpy change when one mole of covalent bonds between atoms in gaseous molecules is broken
(not average when compared to bond dissociation energy)
48
What is the sign of bond dissociation energy?
* Always positive
* Bond dissociation is always endothermic because energy is required to break bonds
49
What does more positive bond energy mean?
The more positive the bond energy, the stronger the bond
50
What is the bond energies for diatomic molecules and for polyatomic molecules?
- bond energy for diatomic molecules are umambiguous
- bond energies for polyatomic molecule are average bond dissociation energies
51
Using bond energy to estimate ΔHᵣ is only applicable to what?
reactions that involve gaseous reactants and products
52
What is the formula for the enthalpy change of reaction?
ΔHᵣ = Σ E(bonds broken) - ΣE(bonds formed)
53
What is standard enthalpy change of neutralisation, ΔHₙₑᵤₜ°?
Enthalpy change when **one mole of water is formed** from the reaction of an acid and an alkali under standard conditions
54
What is the sign of enthalpy change of neutralisation?
* Always negative
* Neutralisation is an exothermic reaction
55
What is the difference between ΔHₙₑᵤₜ° values of strong acid-strong alkali, weak-strong and weak-weak acid alkali?
* Strong acids and strong alkalis isonise completely in dilute aqueous solutions
* Weak-strong alkali acid
* weak acid and alkalis dissociate partially in dilute aqueous solution
* dissociation, which involves bond breaking, is an endothermic process
* during neutralisation, some of the energy released from the neutralisation is absorbed to bring about further dissociation of the weak acid and alkali
* less exothermic, ΔHₙₑᵤₜ° is less negative
* Weak acid-weak alkali
* even less exothermic, ΔHₙₑᵤₜ° is less negative as even more energy is released from the neutralisation is absorbed to bring about complete ionisation of both the weak acid and alkali
56
What is the standard enthalpy change of hydration, ΔH(ₕyd)°?
Standard enthalpy of change of hydration of an ion is the enthalpy change when **one mole of gaseous ions** is hydrated under standard conditions
(from g become aq)
57
What is the sign of ΔH(ₕyd)°?
- ΔH(ₕyd)° is always negative
- hydration process is always exothermic as the hydration process involves formation of ion-dipole interactions between the ions and water molecules
58
What is the magnitude of ΔH(ₕyd)° dependent on?
Charge density of the ion
- with a higher charge and/ or smaller size
- stronger ion-dipole interactions
- more exothermic the hydration reaction
59
What is standard enthalpy change of solution, ΔHₛₒₗₙ°?
Standard enthalpy change of solution of a substance is the enthalpy change when **one mole of the substance is completely dissolved to give an infinitely dilute solution** under standard conditions, so that no further enthalpy change takes place on adding more solvent
60
How does the dissolution of ionic compound occur?
* electrostatic attraction between oppositely charged ions in the solid lattice is overcome to form gaseous ions ⇒ breaking of ionic bonds (energy required is equal to |LE|)
* hydration of gaseous ions ⇒ formation of ion-dipole interactions (hydration energy evolved)
61
What is the formula for standard enthalpy of solution, ΔHₛₒₗₙ°?
ΔHₛₒₗₙ° = -LE + ∑ΔH(ₕyd)°
(note that it's summation)
62
What kind of ΔHₛₒₗₙ° value will the compound dissolve in water?
When ΔHₛₒₗₙ° of an ionic compound is either negative or slightly positive
63
Will ionic compound be soluble in water if the solution process is exothermic, ΔHₛₒₗₙ° < 0?
- The ionic compound will be soluble
- If the solution process is exothermic, it means that the hydration of ions releases more than sufficient energy to break the ionic bonds in the lattice
64
Will ionic compound be soluble in water if the solution process is slightly endothermic, ΔHₛₒₗₙ° slightly > 0?
- compound will be soluble in water
- Although energy released from hydration of the ions is insufficient to break the ionic bonds in the lattice, the compound dissolves as it will result in an increase in entropy of the system
65
Will ionic compound be soluble in water if the solution process is highly endothermic, ΔHₛₒₗₙ° >>> 0?
- compound will be insoluble because energy released from the hydration of ions is insufficient to break the ionic bonds in the lattice
- any increase in entropy of the system is not sufficient to favour its solubility
66
What is the relationship between ΔHₛₒₗₙ° and solubility of the substance in general?
The more negative ΔHₛₒₗₙ° is, the more soluble the substance
67
What is Hess' Law?
Hess' law states that the enthalpy change of a chemical reaction depends only on the initial and final states of the system and is independent of the path taken
68
What is the general formula for ΔHᵣ° using Hess' Law?
ΔHᵣ° = ∑nΔHf°(products) - ∑mΔHf°(reactants)
m and n refer to stoichiometric coefficients of the reactants and products respectively in the balanced equation
69
What is the general formula for ΔHc° using Hess' Law?
ΔHᵣ° = ∑mΔHc°(reactants) - ∑nΔHc°(products)
m and n refer to stoichiometric coefficients of the reactants and products respectively in the balanced equation
70
Why is the experimental values of AgCl, AgBr and AgI more negative than the theoretical values?
- It suggests that the bonding in silver halides is stronger than what the ionic model predicts
- because the bonding is not purely ionic but has substantial degree of covalent character, which is unaccounted for by the ionic model
71
What is a Born-Haber cycle?
An energy cycle showing the various stages in the formation of an ionic compound
72
What does second law of thermodynamics say?
Entropy of the system always increases in a spontaneous process and remains unchanged in an equilibrium process
73
What is entropy, S?
* entropy is the measure of disorder in a system
* the more disordered the system, the higher the entropy
* entropy is the measure of how probable is the state of the system
* the higher the probability of a state occurring, the higher the entropy
74
Why does a disordered state have a higher probability of occurring?
there are more ways of arranging the particles and their energy in a disorder state than a more ordered one
75
What is state function?
Independent on path taken
76
Are enthalpy change, ΔH and entropy change, ΔS state or path functions?
- Both state function
- Both ΔH and ΔS depends only on the initial and final states of the system and independent of the path taken
77
In what processes do entropy increase?
1. Increase in temperature of the substance
2. Change in phase from solid to liquid to gas
3. Reaction leads to increase in number of particles
4. Mixing of particles
5. Gas expansion (increase in volume)
78
How does the change in temperature affect entropy?
- When temperature increases, average kinetic energy of the molecule increases
- there are more quanta of energy available to distribute within the system
- there are more ways to distribute the larger number of quanta at higher temperature
- increase in temperature leads to increase in entropy
78
How does the change in phase affect entropy?
- change in phase: there is a gradual increase in entropy as temperature increases
- When temperature increases, average kinetic energy of the molecule increases
- there are more quanta of energy available to distribute within the system
- there are more ways to distribute the larger number of quanta at higher temperature
- increase in temperature leads to increase in entropy
(basically temperature explanation)
79
What does the graph of entropy against temperature look for phase changes?
- gradual increase in entropy within a phase
- sharp increase with a phase change
79
What is the order of entropy for different phases in increasing order?
* entropy of solid < entropy of liquid < entropy of gas
* Melting: there is an increase in entropy when solid becomes a liquid, volume change is negligible but there is an increase in disorder as the orderly arrangement of particles in the solid is destroyed, there are more ways to arrange the particles and distribute the energy in its liquid state than solid state
* Boiling: there is an even larger increase in entropy when a liquid becomes a gas, gaseous state is the most disordered: particles move randomly in all directions, leading to a large increase in disorder, furthermore, the large increase in volume going from liquid to gas also leads to an increase in entropy
79
How does the change in number of particles affect entropy?
- gas is the most disordered state as they move randomly in all directions
- the formation of greater number of gas particles, entropy of system increases
80
When will gas expand?
If the gas is under pressure, it spontaneously expands when the pressure is released
80
How does gas expansion affect entropy?
Upon gas expansion, there are more ways for gas particles to arrange themselves in a larger volume, leading to an increase in entropy
81
How does the mixing of gas affect entropy?
- Gases always mix completely to create a more disordered state
- each gas expands to occupy the whole container and there are now more ways for the particles to arrange themselves in the larger volume, thus entropy increases
82
What is the effect on entropy when liquids are mixed together?
- liquids of similar polarities mix together towards increased disorder to form a solution in which the particles are uniformly and randomly mixed
- if polarities of liquids are very different, need consider intermolecular forces of attraction
83
What does the spontaneity of reaction depend on?
- enthalpy change
- entropy change
84
What is the Gibbs Equation?
ΔG° = ΔH° -TΔS°
ΔH° = standard enthalpy change of reaction
85
What is the units of ΔH°?
KJ mol⁻¹
86
What is the units of ΔS°?
KJ mol⁻¹ K⁻¹
87
What is the sign of ΔG° used to predict?
spontaneity or feasibility of a reaction or process
88
What does ΔG° < 0 mean?
reaction/process is spontaneous
89
What does ΔG° > 0 mean?
reaction/process is not spontaneous
90
What does ΔG° = 0 mean?
neither the forward nor the reverse reaction/process is favoured; the reaction is at equilibrium
91
What might a reaction not proceed even if ΔG° < 0?
- reaction may be too slow to be observable
- thermodynamics tells us only about the spontaneity of the reaction (whether it can occur) but does not tell us how long it will take to occur (may be high Ea to prevent reaction taking at a reasonable rate)
92
What is the ΔH and ΔS for melting and boiling?
* ΔH > 0: endothermic as intermolecular forces of attractions are overcome
* ΔS > 0: increase in entropy as disorder increases from solid to liquid to gas
93
What is the ΔH and ΔS for decomposition reactions?
* ΔH > 0: most decomposition reactions are endothermic as the bonds in the reactants are stronger than the products
* ΔS > 0: increase in entropy because there is an increase in the amount of gaseous molecules
94
What is the ΔH and ΔS when mixing liquids of similar polarities (eg. hexane and heptane)?
* ΔH ≈ 0: type of imf between molecules in the solutions and in pure liquids is the same
* ΔS > 0: increase in entropy due to mixing of particles
95
What is the ΔH and ΔS when mixing liquids of different polarities (eg. hexane and water)?
* ΔH > 0: endothermic because energy is needed to mix the molecules since in separate liquids, they were held together by different types of imf
* ΔS > 0: increase in entropy due to mixing of particles
96
What is the ΔH and ΔS when dissolving ionic solids in water?
* ΔH > 0 or ΔH < 0: endothermic or exothermic as ΔHₛₒₗₙ° = -LE + ∑ΔH(ₕyd)°
* ΔS > 0: when ionic solids dissolve in water, entropy increases because the ordered ionic lattice is broken up and the ions are then free to move in the solution
97
Why is there different solubilities of ionic compounds in water?
* for solvent particles, entropy decreases because the water molecules that were originally free to move become restricted as they order themselves around the ions during hydration
* the complex interplay between enthalpy and entropy (solute and solvent considerations) cause there to be differing solubilities
98
What is the ΔH and ΔS during condensation or freezing?
* ΔH < 0: exothermic as the intermolecular forces of attraction are formed between the molecules of product
* ΔS < 0: decrease in entropy from a more disordered gaseous/liquid state to more ordered solid state
99
What is the ΔH and ΔS for the decomposition of ozone and dinitrogen monoxide?
* ΔH < 0: exothermic as bonds in the products are stronger than reactants
* ΔS > 0: increase in entropy as there is an increase in the number of gaseous molecules
100
What is the ΔH and ΔS for organic combustion reactions?
* ΔH < 0: Combustion of organic fuels at standard conditions is always exothermic
* ΔS > 0: increase in entropy because there is an increase in the number of gaseous molecules produced
101
What is the ΔH and ΔS for explosions?
* ΔH < 0: exothermic as more energy is released in the formation of bonds in products than the energy absorbed in breaking of bonds in reactants
* ΔS > 0: increase in entropy because there is an increase in the number of gaseous molecules
102
How does an explosion occur?
At high temperatures, reaction rapidly gives off large quantities of gas, resulting in a large increase in entropy, leading to an explosion
103
What reaction has ΔH > 0 and ΔS < 0?
Photosynthesis
104
What are the conditions for spontaneous reaction when ΔH > 0 and ΔS > 0?
* non-spontaneous at low temperatures
* spontaneous at high temperatures
105
What are the conditions for spontaneous reaction when ΔH < 0 and ΔS < 0?
* spontaneous at low temperatures
* non-spontaneous at high temperatures
106
What are the conditions for spontaneous reaction when ΔH < 0 and ΔS > 0?
spontaneous at all temperatues
107
What are the conditions for spontaneous reaction when ΔH > 0 and ΔS < 0?
* non-spontaneous at all temperatures
* has to be driven
108
What are some examples of processes where ΔH > 0 and ΔS > 0?
* melting and boiling point
* decomposition
* some cases of dissolving
109
What are some examples of processes where ΔH < 0 and ΔS < 0?
* condensation and freezing
* addition reactions
* precipitation
* reaction in galvanic cell
110
What are some examples of processes where ΔH < 0 and ΔS > 0?
* some cases of decomposition reaction
* organic combustion reactions
111
What are some examples of processes where ΔH > 0 and ΔS < 0?
Photosynthesis
112
What is ΔG for melting and what are the implications?
* ΔG = 0, ΔH = TΔS
* ΔS(fus) = ΔH(fus) / T(fus)
113
What is ΔG for boiling and what are the implications?
* ΔG = 0, ΔH = TΔS
* ΔS(vap) = ΔH(vap) / T(vap)
114
When is ΔH used instead of ΔG?
* when ΔH is much larger in magnitude than -TΔS
* when both terms act in same direction
115
What are the reactions that occur at room temperature driven by?
Enthalpy driven
116
What are the reactions that only occur at high temperatures driven by?
entropy driven
117
What is sublimation?
solid to gas
118
What is vapourisation?
liquid to gas
119
What is fusion?
solid to liquid
120
How does mixing of solids affect entropy?
mixing of particles in solids (eg. presence of other metal atoms in alloy) also results in an increase in entropy
