Introduction to Thermodynamics

• Thermodynamics originates from the Greek words "thermé" (heat) and "dynamis" (force).
• Rudolf Julius Clausius and William Thomson (Lord Kelvin) formulated the First and Second Laws of Thermodynamics in 1850.
• Thermodynamics is a physical science studying the relationships between heat and other energy forms (mechanical, electrical, chemical).
• Thermodynamics deals with the energy and work of a system, originating in the 18th century with the discovery of steam engines.
• Thermodynamics is concerned with the large-scale responses of a system that can be observed and measured.
• Thermodynamics is a macroscopic science dealing with bulk systems, not the molecular constitution of matter.
• Thomson first used "thermo-dynamic" in 1849 to describe devices converting heat into motion.
• In 1854, Thomson used "thermodynamics" to describe the science of heat engines.
• Thermodynamics studies heat energy and its relation to other energy forms, focusing on initial and final states, not transformation rates.

Branches of Thermodynamics

• Classical thermodynamics analyzes matter macroscopically.
• Classical thermodynamics uses temperature and pressure to calculate properties and predict characteristics.
• Statistical thermodynamics examines every molecule and their interactions to characterize group behavior.
• Chemical thermodynamics studies how work and heat relate in chemical reactions and state changes.
• Equilibrium thermodynamics studies energy and matter transformations approaching equilibrium.

Basic Concepts

• Thermodynamics has a unique vocabulary, where understanding basic concepts is key.
• A system is where energy and/or matter reside, separated from surroundings by boundaries.
• Boundaries are closed surfaces surrounding a system and can be real or imaginary, fixed or deformable.
• Surroundings interact with and influence the system.
• The three types of thermodynamic systems are closed, open, and isolated.
• A closed system (or control mass) has no mass transfer across its boundary, only energy transfer.
• Examples of closed systems are refrigerators and gas compression in piston-cylinder assemblies.
• An open system (or control volume) allows mass flow through the boundary, enabling mass and energy transfer.
• Open systems enclose devices involving mass flow, like compressors, steam turbines, and nozzles.
• An isolated system exchanges neither mass nor energy with surroundings.
• The universe is considered an isolated system.

Properties of a System

• Any characteristic of a system is considered a property.
• State properties define the physical condition of a substance (temperature, pressure, density, specific volume).
• Transport properties measure diffusion within a medium, such as viscosity and thermal conductivity.
• Intensive properties are size-independent (temperature, pressure, density).
• Extensive properties depend on size or extent (mass, volume, total energy).

State Properties - Temperature

• Temperature indicates hotness or coldness.
• Celsius, or Centigrade, is named after astronomer Andres Celsius;
• Fahrenheit is named after German Physicist Gabriel Daniel Fahrenheit.
• Kelvin is named after British Scientist William Thomson (Lord Kelvin).
• Rankine is named after William Macquorn Rankine.
• Absolute temperature is measured from absolute zero.
• Absolute zero is the temperature at which molecules stop moving.
• Absolute zero is equivalent to 0°K (-273.15°C) or 0°R (-460°F).

Temperature Conversion Formulas

• F = (9/5)°C + 32
• C = (5/9)(°F - 32)
• K = °C + 273 = °C + 273.15
• R = °F + 460 = °F + 459.67
• Temperature interval is the difference between two temperature readings on the same scale.

State Properties - Pressure

• Pressure is the force exerted per unit area.
• Gage pressure is measured relative to atmospheric pressure.
• Absolute pressure is measured above a perfect vacuum.

Formulas for Pressure

• Pabs = Pgage + Patm
• Patm at standard conditions is 101.325 kPa, 760 mmHg, 1.013 bar, 14.7 psi, 760 torr, 1.013x10^6 dyne/cm², 1.032 kg/cm², 1 atm, 0 Kpag, 0 psig, or 29.92 in.Hg.
• Negative gage pressure indicates a vacuum.
• Perfect vacuum is 101.325 kPa.
• Critical pressure is the minimum pressure needed to liquefy a gas at its critical temperature.

Heat and Entropy

• Heat is energy associated with the kinetic random motion of many molecules, transferring from hotter to cooler bodies.
• Sensible heat changes temperature without changing phase, Qs = mCΔT.
• Latent heat changes a substance's phase without changing its temperature

Latent Heat Formulas

• QL = ±mL
• Latent Heat of Fusion represents solid to liquid, Latent Heat of Vaporization from liquid to gas.
• Sublimation directly changes solid to gas, while deposition reverses this.
• Evaporation changes liquid to gas; condensation is its reverse.
• Freezing changes liquid to solid, and melting reverses this.
• Entropy is the measure of molecular randomness.

Entropy Rules

• Friction increases entropy.
• Adding heat increases entropy, and removing heat decreases it.
• No entropy change occurs during a reversible adiabatic process.
• For a reversible adiabatic process, ΔS = 0.

Enthalpy and Internal Energy

• Enthalpy is heat energy transferred at constant pressure.
• Internal energy is stored within a body, summing kinetic and potential energies of constituent particles.
• H = U + PV, where H is enthalpy, U is internal energy, P is absolute pressure, and V is volume.

Laws of Thermodynamics

• The First Law is the Law of Conservation of Energy, stating energy is neither created nor destroyed, but transformed.
• Energy can be transferred as heat, work, and with the transfer of matter.
• Plants use photosynthesis to convert sunlight into chemical energy, which humans then convert into kinetic energy.
• The sum of energy entering a system equals the sum of energy leaving it.
• PE1 + KE1 + H1 + Q = PE2 + KE2 + H2 + W is the formula for the sum of energy.
• The Second Law states heat cannot transfer from cold to hot without work input.
• Heat cannot convert 100% into work, requiring engines to operate between hot and cold reservoirs.
• Energy exists at different potential levels, and cannot naturally move from lower to higher potential.
• Heat naturally flows from hot to cold objects (increasing entropy).
• It is impossible for a heat engine to operate in a cycle, receiving heat from a high-temperature body and does an equal amount of work.
• The Third Law states the total entropy of pure substances approaches zero as absolute temperature approaches zero.
• As temperature approaches absolute zero, the entropy of a system approaches a constant minimum.
• The Zeroth Law states if two bodies are in thermal equilibrium with a third, they are in thermal equilibrium with each other.

Ideal Gas

• An ideal gas strictly follows Boyle's and Charles' laws and has a compressibility factor approaching one.
• The Equation State of an Ideal Gas is PV = MRT or PV = nRT.
• Boyle's Law: At constant temperature, volume is inversely proportional to pressure (P1V1 = P2V2) in a confined gas.
• Charles' Law: At constant pressure, volume is directly proportional to temperature in a confined gas.
• Gay-Lussac's Law: At constant volume, pressure is directly proportional to temperature in a confined gas.

General Gas Law

• The General Gas Law combines Charles' and Boyle's Laws.
• The First Law of Thermodynamics states energy can be changed from one form to another, not be created or destroyed. AU = Q + W
• U is the total change in the internal energy of the gas
• Q is the total heat flow of the gas     - When Q is negative, heat is being removed from the system     - When Q is positive, heat is being added to the system
• W is the total work done on or by the gas     - When W is negative, work is being done by the system     - When W is positive, work is being done by the system
• For Monatomic Ideal Gas, translational kinetic energy modes have equations: ΔU = (3/2)nRΔT
• For Diatomic Ideal Gas, translational kinetic energy modes have equations: ΔU = (5/2)nRΔT

The Heat Equation

• Governed by the equation: Q = mcΔT (appropriate for constant volume)
• A thermodynamic process: Is any change a system undergoes from one equilibrium state to another and can be reversible or irreversible.
• Path - Series of states through which a system passes during a process.
• Reversible Processes - Can be reversed without trace on the surroundings.
• Irreversible Process - Proceeds spontaneously in one direction and results in increased molecular disorder.

Processes of Ideal Gas

• Isometric Process: Constant volume process.
• Isobaric Process: Constant pressure process.
• Isothermal Process: Constant temperature process.
• Isentropic Process: Constant entropy process, also known as a reversible adiabatic process.
• Adiabatic Process: No heat flow between system and surroundings.
• Polytropic Process: PVn = C, where n is a constant.