Entropy and the Second Law
5. Entropy in Nature
Learning outcomes
- I can identify entropy changes in everyday situations.
- I can explain entropy in biological systems.
- I can describe entropy in climate and environmental systems.
- I can relate entropy to chemical and physical changes.
- I can evaluate real-world examples using the Second Law.
Entropy in Nature
Entropy helps explain why natural processes have a preferred direction. It applies not only to heat engines and laboratory experiments but also to everyday events, living organisms, chemical reactions, weather, ecosystems, oceans, and the Earth's climate system.
Entropy is represented by:
S
and its SI unit is:
J/K
For an isolated system, the Second Law of Thermodynamics states:
ΔS_total ≥ 0
For a spontaneous irreversible process:
ΔS_total > 0
A useful interpretation is:
natural processes tend to disperse energy and move toward macroscopic states with greater numbers of accessible microscopic arrangements.
Entropy Is Everywhere
Many familiar processes involve entropy changes:
- ice melting
- water evaporating
- perfume diffusing through a room
- hot food cooling
- fuel burning
- metals corroding
- organisms metabolizing food
- weather systems transporting energy
- oceans mixing
- refrigerators cooling food
Some parts of these systems can experience:
decreasing entropy
while the entropy of their surroundings increases.
The important principle is:
the total entropy of an isolated system cannot decrease.
Everyday Example: Cooling Coffee
A cup of hot coffee placed in a cooler room spontaneously cools.
Thermal energy transfers:
hot coffee → cooler surroundings
The coffee loses internal energy and its entropy may decrease.
However, the cooler surroundings receive this energy and gain entropy.
Overall:
ΔS_total > 0
Eventually:
T_coffee ≈ T_room
and thermal equilibrium is approached.
Everyday Example: Ice Melting
Place an ice cube in a warm room.
Thermal energy transfers:
room → ice
The ice changes:
solid → liquid
The molecules become less constrained and have access to more microscopic arrangements.
Therefore:
entropy of the water increases
during melting.
Everyday Example: Water Freezing
Freezing seems more puzzling because:
liquid → solid
The molecules become more ordered.
Therefore:
ΔS_water < 0
Does this violate the Second Law?
No.
Freezing releases thermal energy to the surroundings.
The surroundings gain entropy.
Under conditions where freezing occurs spontaneously:
ΔS_total > 0
So local entropy can decrease while total entropy increases.
Everyday Example: Evaporation
During evaporation:
liquid → gas
Molecules escape from the liquid and gain access to a much larger range of positions and motions.
Therefore:
entropy generally increases
for the evaporating substance.
This is one reason gases generally have greater entropy than liquids.
States of Matter and Entropy
For the same substance under comparable conditions, the general trend is:
solid → liquid → gas
with:
S_solid < S_liquid < S_gas
Solid
Particles are strongly constrained.
Liquid
Particles have greater freedom of movement.
Gas
Particles can occupy vastly more positions and energy states.
Therefore gases generally have the greatest entropy.
Diffusion
Suppose perfume is sprayed into one corner of a room.
Initially:
perfume molecules concentrated in one region
Eventually:
perfume molecules dispersed throughout room
The dispersed state corresponds to enormously more microscopic arrangements.
Therefore:
entropy increases
Diffusion is a spontaneous irreversible process.
Food Colouring in Water
Place one drop of food colouring into still water.
Initially, the dye molecules are concentrated.
Over time:
dye diffuses through water
Eventually the colour becomes approximately uniform.
The reverse process—dye molecules spontaneously gathering back into one drop—is extraordinarily improbable.
This provides a simple visible demonstration of entropy increase.
Mixing
Mixing generally increases entropy because particles gain access to more possible arrangements.
Examples include:
- gases mixing
- salt dissolving in water
- milk mixing with coffee
- pollutants dispersing through air
- substances mixing in oceans
The mixed state usually corresponds to:
more accessible microstates
than the separated state.
Breaking and Shattering
Consider a glass falling and breaking.
Energy initially associated with:
gravitational potential energy
becomes distributed among:
- fragment motion
- sound
- vibration
- deformation
- thermal energy
The total energy is conserved.
But it becomes more widely dispersed.
The fragments do not spontaneously reassemble while collecting all the dispersed sound and thermal energy.
The process is therefore:
irreversible
Friction
A moving bicycle eventually slows if the rider stops pedaling.
Mechanical energy is transformed into internal energy through:
- tire deformation
- bearing friction
- air resistance
- braking friction
Initially:
energy organized as macroscopic motion
Eventually:
energy dispersed among microscopic motions
Therefore:
entropy increases
Entropy in Biological Systems
Living organisms are highly organized.
Cells contain complex structures such as:
- membranes
- proteins
- DNA
- organelles
At first glance, this appears to contradict the Second Law.
It does not.
Living organisms are:
open systems
They exchange both matter and energy with their surroundings.
Organisms Maintain Local Order
Living organisms can create and maintain highly organized structures.
For example:
amino acids → proteins
small molecules → DNA
nutrients → cells and tissues
These processes can decrease entropy locally.
However, organisms require energy to create and maintain this organization.
They release:
- thermal energy
- waste products
- carbon dioxide
- water
- other materials
into their environment.
The resulting entropy increase in the surroundings is sufficient for:
ΔS_total ≥ 0
Metabolism and Entropy
Metabolism includes the chemical reactions occurring within organisms.
Food contains stored chemical energy.
Cells transform this energy into:
- useful chemical energy
- mechanical work
- electrical activity
- molecular synthesis
- thermal energy
Not all energy remains available for useful biological work.
Some becomes dispersed as:
thermal energy
Therefore metabolism contributes to an overall increase in entropy.
Cellular Respiration
During cellular respiration, cells extract energy from molecules such as glucose.
A simplified overall reaction is:
C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy
Some of the released energy is temporarily captured in molecules such as:
ATP
Much is eventually transferred to the environment as thermal energy.
Therefore cellular respiration contributes to:
overall entropy increase
ATP and Useful Energy
Cells use ATP as an important energy-transfer molecule.
Energy from chemical reactions can be coupled to processes such as:
- muscle contraction
- active transport
- protein synthesis
- cell division
However, no biological energy conversion is perfectly efficient.
Some energy becomes:
thermal energy
and is dispersed into the environment.
This is consistent with the Second Law.
Humans and Entropy
A human body maintains a highly organized internal structure.
To do so, it continually:
- consumes food
- takes in oxygen
- carries out chemical reactions
- releases carbon dioxide
- releases waste
- transfers thermal energy to the surroundings
Therefore:
local biological organization does not imply decreasing total entropy
The body maintains organization by continuously exchanging energy and matter with its environment.
Growth and Entropy
Consider a tree growing from a seed.
The tree develops highly organized:
- cells
- tissues
- roots
- leaves
- wood
This appears to create order.
But the tree is not isolated.
It receives:
- sunlight
- water
- carbon dioxide
- minerals
and releases energy and matter into its surroundings.
The total entropy of the larger system can still increase.
Photosynthesis and Entropy
Plants use incoming solar energy to build energy-rich molecules.
A simplified representation is:
6CO₂ + 6H₂O + light energy → C₆H₁₂O₆ + 6O₂
Photosynthesis creates locally organized chemical structures.
But Earth receives high-quality, relatively concentrated solar radiation and ultimately emits energy back toward space as more dispersed infrared radiation.
Therefore biological organization can exist while remaining fully consistent with the Second Law.
Ecosystems and Energy Flow
Energy enters most ecosystems primarily through:
sunlight
Plants capture a fraction of this energy.
Energy then passes through:
producers → consumers → decomposers
At every stage, some energy becomes dispersed as thermal energy.
Therefore energy flow through ecosystems is:
one-directional
Energy is not endlessly recycled in the same useful form.
Matter Cycles, Energy Flows
An important ecological distinction is:
matter cycles
but:
energy flows
Atoms can move repeatedly through:
- carbon cycle
- nitrogen cycle
- water cycle
But useful energy entering from the Sun is gradually dispersed as thermal energy and eventually radiated away.
The Second Law helps explain this difference.
Food Chains
At each trophic level, organisms use energy for:
- movement
- metabolism
- growth
- reproduction
- maintaining body functions
Some energy is transferred to the next trophic level.
Much becomes dispersed as:
thermal energy
Therefore less usable energy is available at higher trophic levels.
The Second Law contributes to understanding why food chains cannot contain unlimited numbers of trophic levels.
Entropy and Decomposition
When an organism dies, decomposers break complex biological molecules into simpler substances.
Energy stored in chemical structures becomes distributed through:
- decomposer metabolism
- thermal energy
- chemical products
Decomposition therefore contributes to the overall dispersal of energy and matter.
Earth Is Not an Isolated System
A critical point when discussing entropy and nature is:
Earth is not energetically isolated.
Earth receives energy from:
the Sun
and emits energy into:
space
This continuous energy flow allows complex structures and processes to develop on Earth while the Second Law remains satisfied.
Solar Energy and Earth
The Sun's surface temperature is roughly:
5800 K
Earth receives solar radiation largely at relatively short wavelengths.
Earth, being much cooler, emits radiation mainly at:
infrared wavelengths
The outgoing energy is more dispersed and associated with lower characteristic temperatures.
This Sun-Earth-space energy flow is fundamental to understanding entropy production in Earth's environment.
Entropy and Climate
Earth's climate system constantly redistributes energy.
Solar heating is uneven because:
- equatorial regions receive more solar energy than polar regions
- day and night alternate
- seasons change solar exposure
- land and oceans heat differently
- clouds and surfaces absorb and reflect differently
These differences create:
temperature gradients
which drive energy transport.
Atmospheric Heat Transport
Warm regions transfer energy toward cooler regions through:
- convection
- atmospheric circulation
- winds
- evaporation and condensation
- radiation
Large-scale atmospheric circulation helps redistribute energy:
warm regions → cooler regions
This reduces temperature differences and produces entropy.
Ocean Circulation
Oceans also redistribute enormous amounts of thermal energy.
Ocean currents transport energy between:
- tropical regions
- temperate regions
- polar regions
These flows are influenced by:
- temperature
- salinity
- density
- wind
- Earth's rotation
Heat transport from warmer to cooler regions is consistent with the Second Law.
The Water Cycle and Entropy
The water cycle involves:
- evaporation
- condensation
- precipitation
- freezing
- melting
- runoff
Solar energy drives much of this cycle.
Evaporation absorbs energy and produces water vapor.
Condensation releases thermal energy into the atmosphere.
These processes transport and disperse energy throughout Earth's climate system.
Hurricanes and Entropy
A hurricane is a highly organized structure.
Does its formation violate the Second Law?
No.
A hurricane is an open system driven by energy transfers.
Warm ocean water supplies energy through:
evaporation
Water vapor rises and condenses, releasing:
latent heat
This drives atmospheric motion while large amounts of energy are ultimately dispersed into the environment.
Local organization can therefore emerge while:
total entropy increases
Weather Can Create Local Organization
Clouds, thunderstorms, hurricanes and other weather structures may appear highly organized.
The Second Law does not say that:
everything must become visually more disorganized
It says total entropy must not decrease in an isolated system.
Energy flowing through a non-equilibrium system can temporarily create organized structures while producing even greater entropy elsewhere.
Chemical Reactions and Entropy
Chemical reactions can involve important entropy changes.
Entropy may increase when:
- solids produce gases
- the number of gas particles increases
- substances mix
- energy becomes more dispersed
Entropy may decrease locally when:
- gases form liquids or solids
- particles become more constrained
- highly ordered structures form
However, spontaneity depends on the total thermodynamic changes, not simply whether the products "look more disordered."
Combustion
Consider burning a hydrocarbon fuel.
Chemical energy is transformed into:
- thermal energy
- radiation
- rapidly moving molecules
- gaseous products
Energy becomes widely dispersed into the surroundings.
Combustion is therefore strongly:
irreversible
The products do not spontaneously gather the dispersed energy and reform the original fuel.
Dissolving
When a substance dissolves:
solute + solvent → solution
particles become distributed throughout the solvent.
For many systems, this increases the number of accessible arrangements.
For example:
salt crystals → ions dispersed through water
However, dissolution involves both energy and entropy effects, so not every substance dissolves spontaneously under every condition.
Chemical Spontaneity
At constant temperature and pressure, chemical spontaneity is often analyzed using:
Gibbs free energy
ΔG = ΔH − TΔS
where:
- ΔG = Gibbs free-energy change
- ΔH = enthalpy change
- T = absolute temperature
- ΔS = entropy change of the system
A process is thermodynamically favorable under these conditions when:
ΔG < 0
This shows that entropy is important, but it is not the only factor determining whether a chemical reaction is spontaneous.
Physical Changes and Entropy
Phase changes provide clear examples.
Melting
solid → liquid
Generally:
ΔS > 0
Vaporization
liquid → gas
Generally:
ΔS > 0
Sublimation
solid → gas
Generally:
ΔS > 0
Freezing
liquid → solid
For the substance:
ΔS < 0
Condensation
gas → liquid
For the substance:
ΔS < 0
Whether each process occurs spontaneously depends on environmental conditions and the total entropy change.
Entropy Change During a Phase Change
For an ideal reversible phase change at constant temperature:
ΔS = Q_rev/T
For melting at the melting temperature:
ΔS_fusion = L_f/T
per unit mass when L_f is the specific latent heat of fusion.
For vaporization:
ΔS_vaporization = L_v/T
per unit mass.
Because vaporization usually produces a major increase in molecular freedom, its entropy increase is often substantial.
Worked Example: Melting Ice
Suppose:
0.20 kg
of ice melts reversibly at:
273 K
and:
L_f = 334,000 J/kg
Energy absorbed:
Q = mL_f
Q = 0.20 × 334,000
Q = 66,800 J
Entropy increase:
ΔS = Q/T
ΔS = 66,800/273
ΔS ≈ 245 J/K
The entropy of the melting ice increases.
Worked Example: Freezing
If the same amount of water freezes reversibly at 273 K:
ΔS_water ≈ −245 J/K
The water loses entropy.
However, the surroundings receive:
66,800 J
of thermal energy.
For a real spontaneous freezing process, the total entropy change of:
water + surroundings
is positive.
Entropy and Corrosion
Rusting is a chemical process involving iron, oxygen and often water.
Once rust has formed, it does not spontaneously transform back into pure iron while restoring oxygen to the environment.
Returning iron oxide to metallic iron requires:
energy and processing
Corrosion therefore illustrates the directionality of real chemical processes.
Human Energy Use
When fuel is burned in:
- cars
- aircraft
- factories
- power stations
concentrated chemical energy becomes distributed through:
- useful work
- thermal energy
- exhaust gases
- sound
- radiation
The total energy is conserved.
But much becomes less available for further useful work.
This is a practical consequence of:
entropy production
Power Stations
A thermal power station converts some thermal energy into:
electrical energy
However, not all the input energy can become electricity.
Some must ultimately be transferred to cooler surroundings.
This is why thermal power stations use:
- cooling towers
- cooling water
- exhaust systems
The Second Law places a fundamental limit on their efficiency.
Refrigerators
A refrigerator creates a region with relatively low thermal entropy by cooling its interior.
But it requires electrical work.
It transfers energy:
cold interior → warmer room
and releases additional energy from the work input.
Therefore:
entropy inside can decrease
while:
total entropy increases
Air Conditioners
Air conditioners operate similarly.
They maintain a cooler indoor environment despite spontaneous heat transfer tending to occur:
hot outdoors → cooler indoors
The air conditioner uses electrical energy to remove thermal energy from the building.
This does not reverse the Second Law because the complete process produces entropy.
Renewable Energy and Entropy
Renewable energy sources also involve entropy.
For example, a solar panel receives relatively concentrated solar radiation and converts some of it into:
electrical energy
The remainder eventually becomes:
thermal energy
No energy conversion is perfectly efficient.
Entropy considerations therefore apply to:
- solar power
- wind power
- hydroelectric power
- geothermal systems
- batteries
The Second Law applies regardless of whether the energy source is renewable.
Waste Heat
Many technologies ultimately release energy as:
waste heat
Examples include:
- engines
- computers
- factories
- power stations
- lighting
- electrical transmission
Calling it "waste" does not mean energy has disappeared.
Instead, the energy has become:
more dispersed and less useful for performing work
Entropy helps explain why recovering all of this energy is impossible.
Entropy and Environmental Systems
Environmental systems contain constant energy and matter flows.
Examples include:
- atmosphere
- oceans
- forests
- rivers
- soils
- ecosystems
These are:
open systems
They can develop local organization because they exchange energy and matter with larger surroundings.
Therefore, environmental order does not contradict the Second Law.
Pollution and Entropy
Imagine a concentrated pollutant entering a lake.
Diffusion, currents and turbulence may spread the pollutant through the water.
The dispersed state is generally more probable.
Returning every pollutant molecule to its original concentrated location requires:
energy and technological intervention
This illustrates an important practical consequence of entropy:
mixing is easy; complete separation can require substantial work
Recycling and Separation
The Second Law does not make recycling impossible.
However, separating mixed materials generally requires:
- energy
- machinery
- sorting
- chemical processing
For example, separating valuable substances from a highly mixed waste stream may require considerably more effort than keeping those materials separated from the beginning.
Entropy therefore has practical implications for:
resource management
Entropy and Climate Change
Entropy should not be confused with climate change itself.
Climate change is primarily studied through Earth's:
- energy balance
- greenhouse effect
- atmospheric composition
- ocean-atmosphere interactions
- carbon cycle
However, all of these processes obey thermodynamic laws.
Greenhouse gases alter how radiation moves through the atmosphere, changing Earth's energy balance and temperature distribution.
The climate system then redistributes this energy through irreversible processes that produce entropy.
Entropy Does Not Mean "Nature Always Becomes Messier"
A common misconception is:
Second Law → everything must become more visually disordered
This is incorrect.
Nature frequently creates organized structures:
- snowflakes
- crystals
- organisms
- hurricanes
- ecosystems
- planetary systems
The Second Law permits local organization.
The requirement is:
total entropy of the complete isolated system does not decrease
Snowflakes
Snowflakes form highly ordered crystalline structures.
During freezing:
water molecules become more ordered
and the entropy of the water decreases.
However, thermal energy is released to the surrounding air.
Therefore:
surroundings gain entropy
The beautiful organization of a snowflake does not violate the Second Law.
Crystals
Crystal formation is another example.
Particles organize into repeating structures.
Locally:
entropy can decrease
But crystal formation often releases energy to the surroundings.
The total entropy change can therefore still be positive.
Order can emerge naturally when the complete thermodynamic balance favors it.
Life Does Not Violate the Second Law
A frequent misconception is that biological evolution or the development of complex organisms contradicts entropy.
They do not.
Earth and living organisms are not isolated systems.
Life receives energy and releases energy into its surroundings.
The Second Law does not prohibit:
local increases in complexity or organization
It requires:
non-decreasing total entropy for the complete isolated system
The Sun-Earth-Space System
A useful large-scale picture is:
hot Sun
↓
relatively concentrated radiation
↓
Earth
↓
biological + atmospheric + geological + technological processes
↓
more dispersed infrared radiation
↓
cold space
This continuous energy flow provides the thermodynamic setting for much of Earth's complex activity.
Entropy and the Direction of Nature
Entropy helps explain why:
hot objects cool
rather than cold surroundings spontaneously heating them further.
Why:
gases mix
rather than spontaneously separate.
Why:
fuel burns
rather than combustion products spontaneously reforming fuel.
Why:
friction slows motion
rather than thermal motion spontaneously accelerating an object.
Why:
organisms require continuous energy
to maintain their structure.
Entropy gives natural processes a preferred macroscopic direction.
The Arrow of Time
Consider watching videos of:
- a glass shattering
- perfume diffusing
- smoke spreading
- ice melting
- hot coffee cooling
You can usually tell immediately whether the video is running forward or backward.
Why?
The forward direction corresponds to:
increasing total entropy
The reverse would require an extremely improbable spontaneous concentration of energy and matter.
This is called the:
thermodynamic arrow of time
Evaluating a Real-World Process
When analyzing entropy in a real system, ask:
What is the system?
What are the surroundings?
Where is energy entering or leaving?
Is matter entering or leaving?
Is energy becoming more concentrated or more dispersed?
Are particles becoming more constrained or less constrained?
Does the system's entropy increase or decrease?
What happens to the surroundings?
Finally ask:
What happens to total entropy?
This avoids the common mistake of examining only one part of a process.
Worked Example: A Tree Growing
Observation
A tree becomes larger and more organized.
Incorrect conclusion
"Entropy decreases, so the Second Law is violated."
Better analysis
The tree is an open system.
It receives:
- solar energy
- water
- CO₂
- minerals
It releases:
- thermal energy
- O₂
- water vapor
- biological waste
The tree can experience local organization while the total entropy of the larger environment increases.
Therefore:
no violation occurs
Worked Example: Melting Ice
System:
ice
Process:
solid → liquid
Particle freedom increases.
Therefore:
ΔS_ice > 0
Thermal energy enters from warmer surroundings.
The process occurs spontaneously when conditions favor melting.
This is consistent with:
ΔS_total > 0
Worked Example: Refrigerator
System of interest:
food inside refrigerator
The food cools.
Its entropy may:
decrease
But the refrigerator uses electrical energy and releases thermal energy to the room.
Therefore:
ΔS_surroundings increases
and:
ΔS_total > 0
No Second Law violation occurs.
Worked Example: Car Braking
A moving car has organized kinetic energy.
Brakes transform much of this energy into:
thermal energy
in:
- brake discs
- tires
- road
- surrounding air
Energy becomes more dispersed.
Therefore:
entropy increases
The process is irreversible.
Worked Example: Pollutant Mixing
A chemical spill enters a river.
Initially:
pollutant concentrated
Later:
pollutant dispersed through large volume
The number of possible arrangements increases dramatically.
Therefore:
entropy of mixing increases
Removing the pollutant requires additional energy and technology.
This demonstrates the practical importance of irreversibility.
First Law and Second Law in Nature
First Law
Energy is conserved.
Second Law
Total entropy of an isolated system cannot decrease.
Therefore natural processes do not destroy energy.
Instead, they tend to:
disperse energy
and reduce the fraction readily available to perform useful work.
Both laws are needed to understand natural systems.
Common Mistakes
Mistake: Saying entropy is simply "messiness."
Entropy is more accurately related to accessible microstates, probability and energy dispersal.
Mistake: Saying entropy must increase in every object.
Local entropy can decrease while total entropy increases.
Mistake: Saying living organisms violate the Second Law because they are organized.
Organisms are open systems exchanging matter and energy with their surroundings.
Mistake: Saying Earth is an isolated system.
Earth continuously receives solar energy and emits radiation to space.
Mistake: Saying freezing violates the Second Law.
The substance loses entropy while its surroundings gain entropy.
Mistake: Saying all phase changes increase entropy.
Melting and vaporization generally increase a substance's entropy; freezing and condensation decrease it.
Mistake: Saying spontaneous means fast.
Spontaneity describes thermodynamic direction, not rate.
Mistake: Saying energy is destroyed when entropy increases.
Energy is conserved but becomes more dispersed.
Mistake: Saying organized structures cannot form naturally.
Local organization can form as long as the complete process obeys the Second Law.
Mistake: Using entropy alone to predict every chemical reaction without considering other thermodynamic quantities.
At constant temperature and pressure, Gibbs free energy provides the appropriate criterion for chemical spontaneity.
Key Terms
Entropy (S) — thermodynamic quantity related to accessible microstates and energy dispersal.
Microstate — one particular microscopic arrangement of particles and energy.
Second Law of Thermodynamics — principle stating that total entropy of an isolated system cannot decrease.
Spontaneous process — process capable of proceeding naturally under the given conditions.
Irreversible process — process that produces entropy and cannot be completely reversed without additional changes.
Open system — system that exchanges energy and matter with its surroundings.
Closed system — system that can exchange energy but not matter with its surroundings.
Isolated system — ideal system that exchanges neither energy nor matter with its surroundings.
Metabolism — collection of chemical reactions occurring within a living organism.
Cellular respiration — biochemical process that releases usable energy from energy-rich molecules.
Energy dispersal — spreading of energy among more particles, states or locations.
Thermal equilibrium — state in which there is no net heat transfer between systems.
Entropy of mixing — entropy change associated with substances becoming mixed.
Gibbs free energy — thermodynamic quantity used to assess spontaneity at constant temperature and pressure.
Thermodynamic arrow of time — preferred direction of macroscopic processes associated with increasing total entropy.
Key Takeaways
- Entropy applies to everyday, biological, chemical, physical, environmental and technological systems.
- For an isolated system, ΔS_total ≥ 0.
- Spontaneous irreversible processes produce entropy.
- Hot objects naturally transfer energy to colder surroundings.
- Diffusion and mixing generally increase entropy.
- Friction disperses organized mechanical energy and increases entropy.
- Melting, vaporization and sublimation generally increase the entropy of a substance.
- Freezing and condensation decrease the entropy of the substance, but can still occur spontaneously because the surroundings gain entropy.
- Local entropy decreases do not violate the Second Law.
- Living organisms are open systems.
- Organisms maintain internal organization by consuming energy and releasing energy and matter to their surroundings.
- Metabolism produces entropy and releases thermal energy.
- Biological complexity does not violate the Second Law.
- Ecosystems receive energy, primarily from the Sun, and ultimately disperse it as thermal radiation.
- Matter can cycle through ecosystems, but useful energy flows through them and becomes progressively dispersed.
- Earth is not an isolated system; it receives solar radiation and emits infrared radiation to space.
- Atmospheric and ocean circulation redistribute energy from warmer toward cooler regions.
- Weather systems can become locally organized while total entropy increases.
- Chemical reactions involve both energy and entropy changes.
- At constant temperature and pressure, ΔG = ΔH − TΔS can be used to assess thermodynamic favorability.
- Combustion is a strongly irreversible process that disperses energy.
- Pollution and mixing illustrate why separation often requires additional work.
- Energy technologies cannot avoid the Second Law.
- Waste heat still contains energy, but that energy is more dispersed and less available for useful work.
- The formation of crystals, snowflakes and living structures is fully compatible with entropy increase.
- The Second Law does not say that everything must become visually disorganized.
- Entropy helps explain the direction of natural processes and the thermodynamic arrow of time.
- Real-world entropy analysis must consider both the system and its surroundings.
Check Your Understanding
- Give three everyday processes that involve an increase in entropy.
- Explain why melting ice generally increases the entropy of the water.
- Why does freezing water not violate the Second Law?
- Explain why diffusion is associated with increasing entropy.
- Why does friction increase entropy even though total energy is conserved?
- Explain why living organisms do not violate the Second Law.
- How can an organism maintain a low-entropy internal structure?
- Explain how cellular respiration contributes to entropy production.
- Why does useful energy become less available as it moves through an ecosystem?
- Explain the statement "matter cycles, but energy flows" in ecosystems.
- Why is Earth not an isolated system?
- Describe how energy moves through the Sun-Earth-space system.
- How do atmospheric and ocean circulation relate to entropy production?
- Explain why the formation of a hurricane does not violate the Second Law.
- Compare the entropy changes associated with melting, freezing, vaporization and condensation.
- Explain how combustion illustrates irreversibility and entropy increase.
- Why can separating a mixed pollutant from water require substantial energy?
- A refrigerator decreases the entropy of its contents. Explain why this does not violate the Second Law.
- Evaluate the statement: "The existence of highly organized living things proves that entropy does not always increase."
- Choose one natural or everyday process and evaluate it using the Second Law by identifying the system, surroundings, energy transfers and overall entropy change.