MCAT Physics · Lesson 3
Thermodynamics
3 min read4 sectionsUpdated
4 sections
3.1 Zeroth Law of Thermodynamics
Introduces thermal equilibrium, temperature scales, and the concept of heat flow between objects.
- Thermodynamics
- Study of energy flow in the universe
- Zeroth Law
- If object A in thermal equilibrium with B, and B with C, then A and C in equilibrium
- Temperature
- No net heat flow between objects in equilibrium
- Heat flows if not in equilibrium
- Proportional to average kinetic energy
- Difference in temperature → heat flow
- Temperature Scales
- Fahrenheit: 32° freeze, 212° boil
- Celsius: 0° freeze, 100° boil
- Kelvin: C + 273, absolute 0 is 0 K, SI unit
- Conversion:
- 98.6°F = 37°C (human homeostasis)
- For ΔF → ΔC, use only 9/5
- Thermal Expansion
- Linear:
- α = coefficient of linear expansion (K⁻¹)
- Volumetric: , β = 3α
- Linear:
3.2 Systems
Defines types of thermodynamic systems and differentiates between state and process functions.
- System Types
- Isolated: no energy/matter exchange, Δtotal = 0 (rare, e.g., universe)
- Closed: energy exchange allowed, no matter exchange (e.g., gas in movable piston)
- Open: energy and matter exchange possible
- State Functions
- Depend only on current equilibrium, not path
- Examples: pressure, density, temperature, volume, enthalpy, internal energy, Gibbs, entropy
- Example: Oil & water separate → ΔS = 0
- Process Functions
- Describe path taken from one state to another
- Include heat (Q) and work (W)
3.3 First Law of Thermodynamics
Relates changes in internal energy to heat and work; covers heat transfer methods and phase changes.
- First Law
- Heat cannot be created or destroyed - only converted from one form to another
-
- +U → internal energy increases, -U → decreases
- +Q → heat into system, -Q → heat out
- +W → work done by system, -W → work done on system
- Heat
- Energy transfer between objects due to temperature difference
- Cannot go from cold → hot spontaneously
- Units: Joule, cal, BTU
- 1 Cal = 1000 cal = 4184 J = 3.97 BTU
- Heat Transfer
- Conduction: via molecular collisions, metals best, gases poorest
- Convection: via fluid motion, liquids & gases only
- Radiation: via electromagnetic waves, works in vacuum
- Specific Heat
- Heat to raise 1 g by 1°C
- Water: 4.184 J/g·K
- Formula:
- Heat of Transformation
- Phase change → temperature constant until complete
- q = mL (latent heat)
- Total heat to melt:
- Fusion: solid ↔ liquid, Vaporization: liquid ↔ gas
- Thermodynamic Processes
- Isothermal: ΔU = 0 → Q = W
- Adiabatic: Q = 0 → ΔU = -W
- Isobaric: constant pressure
- Isochoric: W = 0 → ΔU = Q
3.4 Second Law of Thermodynamics
Explains natural heat flow direction, entropy as energy dispersal, and constraints on reversible/irreversible processes.
- Total entropy of an isolated system always increases.
- Energy Dispersion
- Heat flows from concentrated → spread out
- Entropy (S)
- Measure of energy dispersal
- Formula:
- Units: J/mol·K
- Second Law Implications
- Cold → hot requires work
- Irreversible processes exist; chemical reversibility possible
- Total entropy of universe:
- Example: Ice melting spontaneously in warm environment → natural
- Oil & water separation → reversible → ΔS = 0
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