Module 1 - Thermodynamics - Energy, Entropy, + Spontaneous Change Flashcards
1
Q
What is chemical thermodynamics?
A
Physical chemistry involving… (EPC)
- Energy changes
- Physical processes
- Chemical reactions
2
Q
Why is thermodynamics important in chemistry?
A
- explains energy flow
- transformations (into chemical systems)
- reaction spontaneity = helping predict chemical behaviour
3
Q
What is a physical change?
A
A change where there is no alteration in the chemical composition.
BUT, temperature, pressure, volume, density or phase.. may change
4
Q
What are examples of physical changes?
A
Heating, cooling, expansion, compression, and phase changes.
5
Q
Examples of phase changes? + (how many)
A
6 types
Vaporization, fusion (melting), sublimation, deposition, freezing, condensation
6
Q
Opposite of sublimation?
A
Deposition
7
Q
Opposite of freezing?
A
fusion (melting)
8
Q
Opposite of condensation?
A
Vaporization
9
Q
What does endothermic mean?
A
Energy is required (inputted) = give system high energy = increase KE
10
Q
What does exothermic mean?
A
Energy is released = molecules closer together
= reduce the energy in system = DECREASE KE
11
Q
Example of phases changes that are endothermic?
A
Vaporization (L-G)
Fusion (S-L)
Sublimation (S-G)
12
Q
What is vapourization?
A
Boiling point = (l - g)
- diff bp for diff structures = IMFs
13
Q
What is fusion
A
melting = melting point (s-l)
14
Q
What is sublimation + example
A
s - g
(i.e. dry ice –> gas)
15
Q
Examples of phase changes that are exothermic?
A
- Condensation (g-l)
- freezing (l-s)
- Deposition (g-s)
16
Q
What is condensation + example
A
(g-l)
= boiling water + lid = pasta boiling
17
Q
What is freezing
A
(l-s)
= KE = decreases, molecules slow down
= INTERmolecular force (force BETWEEN) decreases
= Temp decreases
= intermolecular force = closer together = solid
18
Q
What are intermolecular forces
A
Inter = between = interaction bwt. sepreate entitites
hold molecules in a substance together
weaker than INTRA molecular forces
determine the state of matter (s/l/g) + physical properties (i.e. melting/boliling points)
Attractive forces
Dipole,dipole = (polarity)
LDF
Hydrogen bonding (H +… F, O, N)
19
Q
What is intramolecular forces
A
INTRA = between
= forces holding ATOMS in a molecule
STRONGER than intermolecular forces
determine CHEMICAL behaviour of substance
chemical bonds
Covalent, ionic, metal bonds
20
Q
What is a chemical change?
A
A change where the chemical composition of a system = altered by a chemical reaction
21
Q
How do physical and chemical changes differ?
A
Physical changes don’t alter chemical composition, while chemical changes involve a change in the system’s chemical composition due to a reaction.
22
Q
What is deposition? + example
A
(g-s)
- (i.e. frost formation –> H2O vapor (gas) in air = freeze directly into solid (ice crystals))
- (i.e. Snowflakes = H2O vapor (gas) - ice crystals = skip liquid phase)
- Lose energy quickly, KE decreases, = Form SOLID
23
Q
What mechanism does temp. consist of
A
heating/cooling
24
Q
What mechanism does pressure consist of
A
compressing/expanding gas in a specific volume (container)
25
What is density
Density measures how much stuff (mass) is packed into a certain amount of space (volume).
Density = Mass/Volume
(i.e.1 gram of water taking up 1 cubic cm of space, d = 1g/cm^3)
26
How does density relate to physical changes + examples
- either mass or volume of a substance changes without altering its chemical identity.
Heating a Gas:
When a gas is heated, its particles spread out (volume increases), making it less dense.
Freezing Water:
Water becomes less dense as it freezes because ice expands, taking up more space (volume increases).
27
What happens to the chemical composition of a substance during a physical change?
It remains unchanged, but the state or form of the substance might change.
28
Is energy involved in physical changes?
Yes, energy is exchanged during physical changes.
29
In what forms is energy transferred during a physical change?
Energy is often transferred as heat or work.
30
What law of thermodynamics is related to energy transfer during physical changes?
The first law of thermodynamics.
31
What is the first law of thermodynamics?
Mass is conserved.
= Mass cannot be created or destroyed only transformed or transferred.
Energy is conserved
32
What is the second law of thermodynamics?
Entropy does not decrease.
- Depends on how ordered the system is.
Inc. Entropy = Steam
Dec. Entropy = Ice
33
What is a system? + examples
A region of space we are interested in (e.g., engine, chemical reaction vessel).
34
What are surroundings?
Everything outside the system that can interact with it (e.g., to measure heat or work).
35
Open System + example
Both energy and matter can be exchanged with surroundings (e.g., a mug of coffee
36
Closed System
Only energy can be exchanged; matter is contained (e.g., sealed soda can)
37
Isolated System + example
Neither energy nor matter is exchanged (e.g., a perfect thermos)
True isolated systems are theoretical since even thermoses lose heat over time.
- Aim for thermal equilibrium internally because energy stays within the boundaries. However, a true isolated system doesn’t interact with its surroundings to achieve thermal equilibrium externally.
38
Why is a coffee mug not a closed system?
A mug of coffee losing heat and evaporating water.
= losing matter + energy
39
Can a soda can eventually become an open system? When?
Sealed soda can is never an open system because both energy and matter cannot freely exchange with the surroundings.
However, when you open the soda can:
- Matter (like CO₂ gas and possibly soda droplets) escapes. +
- Energy (in the form of sound or heat) is released.
T
his transforms the soda can into an open system because both matter and energy can now move between the system (soda) and surroundings.
Sealed soda can = Closed system (only energy can move in or out).
Opened soda can = Open system (both energy and matter can move in or out).
40
What is "heat" vs. "energy"?
Heat is a type of energy transfer. Specifically, it refers to the movement of thermal energy between the system and surroundings due to a temperature difference.
A warm soda can transfers heat (thermal energy) to the cooler air around it.
Heat is a form of energy transfer caused by a temperature difference, while energy can also refer to other forms like sound, work, or kinetic energy.
41
What are the specific forms of energy (transfer)?
Work: When gas escapes a soda can, the movement of air molecules against atmospheric pressure is a type of work done by the system.
Sound energy: The “psst” sound when opening a soda can is energy released as sound waves.
Thermal energy that flows between the soda (system) and its surroundings.
42
What happens to heat in a cold soda v.s. warm soda?
A cold soda absorbs heat from the warmer room air, becoming less cold over time.
A warm soda can releases heat to the cooler surroundings, losing thermal energy.
43
Is heat used for potential energy or kinetic energy?
Hint: what is the molecular explanation for heat transfer in a pure substance
Temperature increases:
Heat is used to increase the kinetic energy of molecules (their motion). This happens when a substance is heated and the temperature rises without a phase change.
Phase changes:
During a phase change (e.g., melting or boiling), the temperature stays constant, and heat is used to overcome intermolecular forces, which increases potential energy.
44
Example of phase change using molecular explanation
When boiling water, heat breaks the hydrogen bonds between water molecules, allowing them to move apart and transition from liquid to gas.
45
What happens when attractive forces are overcome?
Overcoming attractive forces allows molecules to move further apart, changing the phase:
From solid to liquid: Molecules break out of their rigid lattice and flow.
From liquid to gas: Molecules break free completely and spread out.
From gas to liquid or liquid to solid (aggregation): Molecules come closer together, forming organized, compact structures.
46
Heat (q): Definition
Heat is energy transferred between a system and its surroundings due to a temperature difference.
47
Heat (q) Directionality
Heat flows from high temperature to low temperature until thermal equilibrium is reached.
Positive Q: Heat absorbed by the system (endothermic).
Negative Q: Heat released by the system (exothermic).
48
Thermal Equilibrium
The system and surroundings have the same temperature, so no heat flows (Q = 0).
49
Heat Transfer Formula
and what does it represent
Formula:
𝑄 =𝑚𝑐Δ𝑇
Q: Heat energy (J).
m: Mass (g).
c: Specific heat capacity (J/g°C).
ΔT: Temperature change (°C).
Shows what heat is dependent on:
- Quantity of substance (mass or moles).
- Specific heat capacity.
- Magnitude of temperature change.
50
What happens during a phase change? + when
During a phase change:
Temperature is constant.
Heat is used to overcome intermolecular forces (increasing potential energy).
Examples:
During "transition temp."
Fusion (solid → liquid).
Vaporization (liquid → gas).
51
What happens during no phase change and when does it occur?
= just heating
Temperature rises.
Heat increases kinetic energy of molecules
= during temp change
= slant
52
Types of Boundaries + explanation
Diathermic boundary:
Allows heat transfer between system and surroundings.
Example: Stainless steel container.
Adiabatic boundary:
Does not allow heat transfer (Q=0). between system and surroundings.
Example: Thermos with vacuum-insulated walls.
53
Work and Heat Relationship + ex.
Work: Energy transfer involving force and displacement.
Example: Expansion of gas in a piston.
Heat into the system
(Q>0): Endothermic (system gains energy).
Heat out of the system (Q<0): Exothermic (system loses energy).
54
What does a heating curve show? what does each axis show + curve + slant
A heating curve shows the temperature of a substance as heat is continuously added at constant pressure.
X-axis: Heat added
Y-axis: Temperature
The curve has slants (temperature changes) and plateaus (phase changes).
55
What are the stages of the curve - (slants + plateaus)
1. Solid (Q solid):
Temperature increases as heat is added (slant).
Formula: Q=mCsolidΔT
2. Fusion (Q fusion): Temperature remains constant (plateau) as the solid melts into a liquid.
Formula: Q=mΔHfusion
3. Liquid (Q liquid): Temperature increases again (slant).
Formula:
Q=mC liquidΔT
4. Vaporization (Q vapor): Temperature remains constant as the liquid boils into gas (plateau).
Formula: Q=mΔHvaporization
Gas (Q gas): Temperature increases further (slant).
Formula: Q=mCgasΔT
56
What are the key observations of the heating curve?
Phase changes occur at constant temperatures:
Tfusion = Melting point.
Tvaporization = Boiling point.
57
Why is the qvap greater than qfus?
Heat of vaporization (Q vap) is greater than heat of fusion (Q fusion) because breaking intermolecular forces during vaporization requires more energy.
58
Heat Capacity
Heat capacity measures how much heat is required to raise the temperature of a substance by
1∘ C or 1K
Formula: C= Q/ΔT
Rearranged: Q=CΔT
59
Specific Heat Capacity (C)
Refers to heat capacity per unit mass or mole:
C = Q solid/ mΔT
solid
Units:
J/g∘C
or J/mol∘C
Example: Water has a higher heat capacity than metals because its complex structure can absorb more energy
60
Heat Capacity at Constant Volume vs. Constant Pressure
- which is usually larger + why?
Cv: Heat capacity at constant volume.
𝐶𝑝: Heat capacity at constant pressure (usually larger due to work done by the system).
61
Why Does Heat Capacity Depend on Temperature?
- At higher temperatures, molecules vibrate more energetically and may activate additional degrees of freedom (e.g., rotation, vibration).
- Complex molecules have more ways to store energy, increasing their heat capacity.
62
Key Takeaways for Problems with heat capacity
Match units: Ensure all values are in
K or ∘C and J or kJ
Use the correct formula based on the phase:
Slants: Q=mCΔT.
Plateaus: Q=mΔH.
- Be mindful of whether the process occurs at constant volume or pressure.
63
Relationship 1:
More complex molecules →
More complex molecules → Higher heat capacity.
64
Relationship 2:
Phase changes → Constant ____ → Energy breaks ____ ______ .
Phase changes → Constant temperature → Energy breaks intermolecular forces.
65
Relationship 3:
Heating slants → Energy ____ temperature → Depends on _____
Heating slants → Energy raises temperature → Depends on heat capacity.
66
What is a Pure Substance?
s made up of only one type of particle (atoms or molecules) and has a uniform composition and properties throughout.
Examples:
Elements: Oxygen (O2), Gold (Au)
Compounds: Water, Sodium chloride (NaCl)
67
What is Not a Pure Substance? + ex.
Substances that are not pure include mixtures, where more than one type of particle is present.
Examples:
Homogeneous Mixtures
(Solutions): Saltwater, air.
Uniform composition but made of different particles
Heterogeneous Mixtures: Sand and water, salad.
Non-uniform composition.
68
Why Does Heat Capacity Depend on Purity?
For pure substances, the heat capacity is defined as a constant value because their composition is uniform. This allows us to calculate heat using simplified formulas:
Q=mCΔT
C= specific heat capacity (J/g°C or J/kg°C).
Q=nCmolarΔT
n= moles of the substance.
C molar = molar heat capacity (J/mol°C).
For mixtures, heat capacity can vary depending on the proportion of different substances, making it harder to define a single value.
69
What to do in a heat transfer problem? (ex.1-1)
When mixing two different types of substances at different temperatures to find the final temperature, this is a heat transfer problem. The principle used is energy conservation.
Energy conservation: Energy is neither created nor destroyed, but transformed or transferred between systems, maintaining equilibrium.
The equation used is:
Q=mCΔT
Key Principle: Heat lost by one substance equals heat gained by the other.
Equation:
Q copper =−Q H2O
Make sure to properly account for signs (plus and minus) when calculating.
At constant pressure, no work is done on or by the system because there is no change in volume (solids do not expand or compress significantly).
- In most heat transfer problems involving solids, the volume change is negligible. = the system does not perform work on its surroundings or have work done on it
==> because the volume of solids doesn’t change significantly when heat is added or removed. This simplifies calculations and is why work is not included in the equation.
70
Work + in Chemistry
Work (w) is the product of force and displacement in a system, representing energy transfer through movement.
In chemical systems, work is often associated with the movement or change in position of particles (atoms or molecules)
71
Types of Work
Expansion Work (Compression/
Expansion(heat) of Gases):
This type of work occurs when a gas expands or compresses, causing movement, such as when a gas pushes a piston to lift a mass.
- Expansion work: Gas does work on the surroundings (pushes piston outward).
- Compression work: Surroundings do work on the gas (pushes piston inward).
Electrical Work:
Work can be done by moving charged particles (like electrons) between regions of high and low electrical potential, which is important in processes like ion flow or electrical impulses in neurons
72
Expansion Work in atmosphere
Atmospheric Pressure: The atmosphere exerts a continuous downward pressure on systems.
73
Expansion Work - Work Done by Surroundings
If the surroundings exert pressure that decreases the system's volume, the surroundings do work on the system.
74
Expansion Work - Work Done by Surroundings ----equation
Work Equation: When volume changes (ΔV), and external pressure (Pₑₓ) is constant, the work done by surroundings is:
W = −PₑₓΔV
Negative Sign: The negative sign indicates that work is done on the system when it is compressed (volume decreases) and by the system when it expands (volume increases).
Energy (Work): Measured in joules (J).
75
Work and Thermodynamics (positive and negative meaning)
Positive Work: Work is done on the system by the surroundings when the system is compressed. This typically results in heat absorption, which can be endothermic (if temperature increases).
Negative Work: Work is done by the system on the surroundings when the system expands. This often involves heat release, which can be exothermic (e.g., when gas expands, pushing a piston).
76
exothermic related to work
Exothermic Reaction: Heat is released, and work is done by the system (e.g., expansion of gas generating an explosion).
77
endothermic related to work
Endothermic Reaction: Heat is absorbed, and work is done on the system (e.g., compression of gas).
78
What you can use the work equation for in conjunction to endo/exo
W = −PₑₓΔV
ΔV, you can calculate work done during compression or expansion processes, and by knowing the system’s behavior (whether it absorbs or releases heat),
= predictions
you can determine if a reaction is exothermic or endothermic.
79
First Law + eq.
Energy cannot be created or destroyed, only transformed and transferred.
ΔU=Q+W
ΔU: Change in internal energy.
Q: Heat transferred (+Q: heat added,
(−Q: heat removed).
W: Work
(+W: work on the system,
(−W: work by the system).
In short, work is energy transfer, and heat is energy transformation, but both contribute to changes in internal energy (ΔU).
80
What is transformed energy?
Transformed Energy (Q, Heat):
Energy changes form, typically from one type of molecular motion to another (e.g., thermal energy).
Example: When you heat a gas, the energy added increases the random motion of particles.
81
What is transferred energy?
Transferred Energy (w,Work):
Energy is moved from one part of the system (or surroundings) to another without changing its form.
Example: A gas expands, transferring energy to move a piston (mechanical energy).
82
In thermodynamics what is manifested in organized motion?
Work focuses on organized motion (e.g., all gas particles push outward to expand the system).
83
In thermodynamics what is manifested in random motion?
Heat involves random motion (e.g., faster random movement of particles due to thermal energy).
84
What does internal energy include?
PE + KE
85
What is the first law of thermodynamics?
The first law of thermodynamics states that the change in internal energy (ΔU) of a system is equal to the heat (Q) added to the system plus the work (W) done by the system:
ΔU=Q+W
86
What happens in a constant volume process?
In a constant volume process, the volume of the system does not change, and therefore no work is done by the system (W=0). The heat transferred to the system (qV) equals the change in internal energy (ΔU):
qV=ΔU
87
What is the relationship between heat transferred and internal energy in a constant pressure reaction?
In a constant pressure reaction, heat transferred (qp) is equal to the change in enthalpy (ΔH):
qp=ΔH
This is because at constant pressure, heat transfer is linked to both internal energy and the work done by the system due to volume change.
88
How does work factor into a constant pressure reaction?
In a constant pressure reaction, the volume of the system can change, and this change in volume does work (W=−Pext ΔV) on the surroundings. The heat transferred in this process is given by the change in enthalpy (qP=ΔH=ΔU+PΔV).
89
Why does work not impact the internal energy in a constant volume reaction?
In a constant volume reaction, there is no volume change (ΔV=0), so no work is done by the system (W=0). As a result, the heat transferred to the system goes entirely into changing the internal energy
(qV =ΔU), and pressure does not affect the internal energy.
90
How do you get from the first law to the alternative statements for constant volume and constant pressure?
For constant volume, since no work is done (W=0), the heat transferred equals the change in internal energy:
qv=ΔU
For constant pressure, heat transferred is equal to the change in enthalpy (qp=ΔH), where enthalpy is defined
H=U+PV
91
How is enthalpy (H) defined?
Enthalpy (H) is defined as the sum of the internal energy (U) and the product of pressure (P) and volume (V):
H=U+PV
92
What is the physical interpretation of
ΔH in a constant pressure process?
In a constant pressure process, ΔH represents the heat transferred to or from the system. Since the pressure is constant, the heat change corresponds directly to the enthalpy change.
93
What is the key difference between constant volume and constant pressure processes in terms of energy transfer?
In constant volume, the heat transferred directly changes the internal energy (qv=ΔU), and no work is done.
In constant pressure, the heat transferred changes both the internal energy and the work done due to volume change (qp=ΔH=ΔU+PΔV)
94
Why does pressure have no effect on the internal energy in a constant volume reaction?
anything multiplied by 0 = 0
In a constant volume reaction, pressure does not affect the internal energy directly because no work is done (W=0). The heat added to the system goes solely into changing the internal energy (qv=ΔU).
95
Constant Volume look for phrases?
Constant Volume (Use qv=ΔU):
Look for phrases like:
"Rigid container"
"Sealed vessel"
"No change in volume" (ΔV=0)
In this case, you know no work is done (W=0), and heat transfer directly affects internal energy (ΔU).
96
Constant Pressure look for phrases?
Constant Pressure (Use qp=ΔH):
Look for phrases like:
"Atmospheric pressure"
"Open container"
"Pressure is constant"
In this case, heat transfer affects the enthalpy (ΔH), and the system may perform work
(W=−Pext ΔV).
97
What assumption to make? hint alternative statement
If the problem says "a reaction is carried out at constant pressure, releasing 200 kJ of heat," you can directly conclude:
qp=ΔH=−200kJ
98
What assumption to make?
If it says "a rigid container absorbs 50 J of heat," you know:
qv=ΔU=50J
99
What does the first law of thermodynamics require?
The first law of thermodynamics requires the energy to be conserved, but it places no restriction on how energy is transferred
100
Why are certain processes, like a ball bouncing higher after absorbing energy from the floor or hot copper becoming hotter from contact with cold zinc, not likely to occur in reality?
These processes would violate the second law of thermodynamics because they would lead to a decrease in entropy, which doesn't naturally occur in isolated systems.
101
What is entropy?
Entropy is a measure of the number of possible microstates (configurations) within a system. It connects macroscopic thermodynamic properties like temperature and pressure with microscopic behaviors like the motion of molecules.
102
What are microstates in the context of entropy?
Microstates refer to the different possible microscopic configurations (position, velocity, and energy) of a system, such as the motion of molecules in a gas
103
How does the macroscopic state of a gas relate to its microscopic states?
The macroscopic state of a gas (characterized by temperature, pressure, volume, and moles) is only one of many possible microscopic configurations, where the molecules have random motion and continuous changes in position, velocity, and energy.
104
What is the connection between the macroscopic and microscopic description of a system?
The macroscopic description (e.g., pressure, volume, temperature) simplifies the system, while the microscopic description includes all the individual positions, velocities, and energies of the molecules, which are constantly changing.
105
What principle does the macroscopic description of a system not follow strictly?
The macroscopic description does not strictly follow quantum mechanics, because according to Heisenberg's uncertainty principle, the exact position and velocity of a particle cannot be known simultaneously
106
What does the Boltzmann equation describe?
The Boltzmann equation relates the number of microstates W to the entropy S of a system. It shows that the greater the number of microstates, the higher the entropy.
107
What is the Boltzmann constant
kB related to?
The Boltzmann constant
k is related to the gas constant
R and Avogadro's number NA
by the equation:
kB = R/NA
107
What is the Boltzmann equation for entropy?
The Boltzmann equation is:
S=kBln(W)
S is entropy
𝑘𝐵 is the Boltzmann constant
W is the number of microstates.
108
What is the relationship between the number of microstates and entropy?
The greater the number of microstates (W), the higher the entropy (S), because more configurations lead to more disorder in the system less order. = more possible ways the molecules = can be configured/ordered.
109
How does entropy help in determining the energy distribution of molecules?
Entropy helps in describing how molecules are distributed among the available energy levels, which influences the flow of energy and the state of the system (whether it is more ordered or disordered).
110
Why is entropy important in chemistry?
Entropy is important because it reflects the distribution of energy within a system and predicts the spontaneity of processes. Spontaneous processes tend to increase entropy, which corresponds to energy spreading out or becoming more evenly distributed. This helps us understand why systems naturally evolve toward equilibrium, such as when energy flows from hot to cold or when systems undergo phase transitions.
Entropy helps us understand why certain processes naturally move toward higher entropy states (more disorder). For example, a ball bouncing higher and higher seems impossible in real life because it would require an increase in energy without any external input. Without that input, the system wouldn't reach equilibrium and would violate entropy principles. In a natural system, energy tends to spread out or dissipate, increasing entropy, which is why spontaneous processes go in the direction that maximizes disorder.
111
How does an increase in the number of accessible energy levels relate to microstates and entropy?
An increase in the number of accessible energy levels results in a greater number of possible microstates. This means energy can be dispersed in more ways among those levels, leading to an increase in entropy.
112
Consider a system with 10 units of energy and 2 accessible energy levels. What would happen if more energy levels are present?
10 units of energy and 5 energy levels. In the latter case, energy can be distributed in more ways, leading to higher entropy.
113
Why do translational energy levels contribute the most to entropy?
Translational energy levels are very closely spaced, allowing molecules to access a large number of energy states, leading to more microstates (W) and higher entropy (S).
114
How do temperature and volume affect translational motion?
Temperature: Higher temperatures increase the energy of molecules, allowing access to more translational energy levels, increasing microstates and entropy.
Volume: A larger volume provides more space for molecules to move, increasing the number of accessible translational energy levels and microstates.
115
Why do rotational and vibrational motions contribute less to entropy?
Rotational and vibrational energy levels are more widely spaced than translational levels, so fewer energy states are accessible, leading to fewer microstates and a smaller contribution to entropy.
116
How does temperature influence rotational and vibrational motions?
Rotational: At lower temperatures, fewer rotational energy levels are accessed. At higher temperatures, more levels become accessible, but the contribution is still less than translational motion.
Vibrational: At lower temperatures, molecules remain in their ground vibrational state, contributing minimally to entropy. At higher temperatures, vibrational contributions increase as molecules access higher energy levels
117
Why is translational motion dominant even at higher temperatures?
Translational motion always dominates because of its vast number of closely spaced energy levels. Even at high temperatures, the contribution from rotational and vibrational motions remains smaller in comparison.
118
What does the Second Law of Thermodynamics state?
The entropy of the universe can never decrease.
ΔS universe =ΔS system +ΔS surroundings ≥0
119
What does
ΔS
universe =0 mean?
The process is thermodynamically reversible.
The system and surroundings can return to their original states if the process is reversed.
120
What does
ΔS
universe >0 mean?
The process is spontaneous in the forward direction.
Reversing the process does not restore the universe to its original state.
121
What does
ΔS
universe <0 mean?
The process is impossible or forbidden.
Such a change would violate the second law of thermodynamics.
122
Why can't
ΔS
system
or
ΔS
surroundings
alone predict spontaneity?
The direction of change can only be predicted by considering the total entropy change of the universe,
ΔS
universe
123
What is entropy (S) on a molecular level?
Entropy measures the number of ways a system can be microscopically different due to the distribution of energy among molecules (microstates).
124
Why is
ΔS
universe
important for predicting spontaneity?
If
ΔS
universe>0: The process is spontaneous.
ΔS
universe =0: The process is reversible.
ΔS
universe<0: The process is impossible.
125
Why is entropy always increasing in the universe?
Entropy naturally increases because energy disperses and systems tend to favor more microstates, which align with the second law of thermodynamics.
126
What is an example of a thermodynamically irreversible process?
Cooking an egg: once cooked, the egg cannot return to its raw state.
127
Entropy Changes for the Surroundings
Determined by how much heat (Q) is absorbed or released by the system.
Formula:
ΔS surr = − ( q/Tsurr)
128
Entropy Changes for the Surroundings
For constant pressure processes:
qp =ΔH
So:
ΔS surr = − (ΔH/Tsurr)
129
Sign of ΔS_surr
exo + endo
Exothermic process (heat released by the system):
Q<0, surroundings absorb heat.
ΔS surr >0 (positive).
Endothermic process (heat absorbed by the system):
Q>0, surroundings lose heat.
ΔSsurr <0 (negative).
130
Entropy Change for the System (ΔS_sys)
Entropy depends on the type of process:
should know formula if not check pg 17
- Phase Change (ΔS_sys)
Only valid if the phase change occurs at the transition temperature (e.g., melting, vaporization).
- Constant Volume Heating/Cooling:
Cv : Heat capacity at constant volume
- Constant Pressure Heating/Cooling:
Cp: Heat capacity at constant pressure.
- Change in State for an Ideal Gas:
Combines temperature and volume/pressure effects:
131
What are some notes to keep in mind when applying formulas to the entropy change for the system?
Heat Capacity (C):
Assumes C is constant between
𝑇𝑖 and 𝑇𝑓 (minor fluctuations in heat capacity are neglected).
Ideal Gas Relationship:
For ideal gases,
CP =CV +nR.
Units of Entropy:
Always use Kelvin for temperature.
Units:
J/K
Ttr - represents the transition temperature—the temperature at which a substance undergoes a phase change, such as melting or vaporization.
So, for example:
Fusion is the phase change from solid to liquid (e.g., melting), and its transition temperature would be the melting point.
Vaporization is the phase change from liquid to gas, and its transition temperature would be the boiling point.
The equation used to calculate the entropy change during phase transitions is:
ΔS= ΔH/Ttr
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What happens to entropy when the temperature of a system increases?
Entropy increases because the system has access to more energy levels, which increases the number of microstates.
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How does an increase in temperature affect the number of accessible energy levels?
As temperature increases, more energy levels become accessible, which increases the number of microstates and, therefore, the entropy.
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What is the relationship between temperature, energy levels, microstates, and entropy?
As temperature increases, more energy levels become accessible, leading to an increase in the number of microstates (W), which results in an increase in entropy (S).
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How does increasing the volume of a gas affect entropy?
Increasing the volume of a gas increases the number of accessible translational energy levels, which increases entropy.
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What happens to translational energy levels as volume increases?
The number of accessible translational energy levels increases because the energy levels become more spaced out as the volume expands.
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How does vaporization, sublimation, and fusion affect entropy?
Vaporization and sublimation significantly increase entropy due to the dispersion of energy among more translational energy levels. Fusion also increases entropy but to a lesser extent.
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Why does fusion increase entropy less than vaporization or sublimation?
In fusion, molecules gain some rotational and translational energy as they transition from solid to liquid, but the increase in entropy is smaller compared to vaporization or sublimation, where more energy levels are involved.
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What happens when a solid melts?
When a solid melts, the molecules gain some rotational and translational motion, causing the energy of the system to become dispersed among a greater number of energy levels, which increases entropy.
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Define enthalpy (ΔH) and entropy (ΔS)
Enthalpy (ΔH): Heat content of a system, measures heat absorbed or released at constant pressure.
Entropy (ΔS): Measure of randomness or number of microstates in a system.
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What does it mean when ΔH is negative? What about ΔS being positive?
ΔH<0: Exothermic process, heat released.
ΔS>0: Increase in disorder or number of microstates.
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Why can entropy (ΔS) alone not determine spontaneity?
Entropy doesn’t account for heat transfer (ΔH). Both entropy and enthalpy are needed to calculate Gibbs free energy (ΔG) for spontaneity.
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Write the equation for Gibbs free energy.
ΔG=ΔH−TΔS, where:
ΔH: Enthalpy change (heat content).
T: Temperature in Kelvin.
ΔS: Entropy change (disorder/microstates).
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Under what conditions is a reaction always spontaneous?
When
ΔH<0 (exothermic) and
ΔS>0 (increasing entropy).
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What happens to entropy (ΔS) of the surroundings when heat is released by a system?
Entropy of surroundings increases because the heat transfer increases the number of microstates in the surroundings.
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Why does high temperature magnify the impact of
ΔS in the Gibbs equation?
ΔG=ΔH−TΔS: At high T
TΔS term becomes larger, making entropy (ΔS) dominate the equation.
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Why is freezing water an exothermic process?
Freezing releases heat (ΔH<0) because molecules lose energy as they form a more ordered solid structure.
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When is a process thermodynamically reversible?
When ΔS universe=0
and
ΔG=0.
148
Explain why ice melts spontaneously at room temperature but not in a freezer.
At room temperature:
ΔG=ΔH−TΔS<0, so melting is spontaneous.
In a freezer:
T is too low for
TΔS to overcome
ΔH, making
ΔG>0.
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Why is condensing water vapor spontaneous at low temperatures but not at high temperatures?
At low
T: Δ
ΔH<0 dominates because
TΔS is small, making
ΔG<0
At high TΔS > ΔH, making ΔG>0.
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Can a process with
ΔS<0 ever be spontaneous?
Yes, if
ΔH<0 and large enough to make
ΔG<0, even with decreasing entropy.
151
What do
ΔH,
ΔS,
ΔG, and
ΔU represent in thermodynamics?
They represent changes in enthalpy, entropy, Gibbs free energy, and internal energy, respectively, often calculated for a reaction.
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What does the subscript "r" in
indicate?
It denotes these quantities are calculated for a reaction.
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What is
ΔHr (enthalpy of reaction)?
It's the enthalpy change per mole of reaction, measured in joules per mole. It's an intensive quantity, meaning it does not depend on the system size.
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What is an extensive quantity?
A property that depends on the size of the system (e.g., ΔH).
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What is an intensive quantity?
A property that does not depend on the system size, like ΔHr or density.
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Why is density an example of an intensive quantity?
The density of water remains the same whether you measure 1 mL or 1 L because it's a ratio of two extensive quantities (mass and volume).
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