Introduction to Thermodynamics

Thermodynamics Definition

• Thermodynamics studies energy and its impact on substances' physical properties.
• The term "thermodynamics" has Greek origins stemming from "therme" (heat) and "dynamics" (power).
• Thermodynamics is the field of science focused on the correlation of heat, work, and system properties while in equilibrium.

Applications of Thermodynamics

• Examples of thermodynamics at work include the human body, air conditioning, airplanes, car radiators, power plants, refrigeration systems, solar collectors, showers, hot water, cold water, heat exchangers, and pumps.

Units and Dimensions

• Here are the multiples and prefixes used in science:
  • 10^12 is Tera, symbol T
  • 10^9 is Giga, symbol G
  • 10^6 is Mega, symbol M
  • 10^3 is Kilo, symbol k
  • 10^-3 is milli, symbol m
  • 10^-6 is micro
  • 10^-9 is nano, symbol n
  • 10^-12 is pico, symbol p
• Dimensions, units, and symbols:
  • Length is measured in meters (m).
  • Mass is measured in kilograms (kg).
  • Time is measured in seconds (s).
  • Temperature is measured in Kelvin (K).
  • Electric current is measured in Ampere (A).
  • Amount of light is measured in Candela (cd).
  • Amount of matter is measured in moles (mol).

Temperature

• Temperature gauges the degree of hotness or coldness of a body.
• 30°C is equivalent to 303 K (Kelvin).

Specific Volume

• Specific volume equals volume divided by mass.
• The unit for specific volume is m³/kg.

Pressure

• 1 Pascal (Pa) equals 1 N/m².
• Standard atmosphere is 1 atm, which translates to 101.325 kPa or 1.01325 bar.
• 1 bar equates to 10^5 Pa, 100 kPa, or 0.1 MPa.

Density

• Density (ρ) represents the mass of a substance per unit volume.
• Density is measured in kg/m³.
• Density equals mass divided by volume (ρ = m/V).

Macroscopic vs. Microscopic Approaches

• Macroscopic Approach (Classical Thermodynamics):
• Focuses on the overall quantity of matter.
• Doesn't consider molecular-level events.
• Studies thermodynamics without requiring knowledge of individual particle behavior.
• Deals with the effects of numerous molecules, detectable by human senses.
• Observations are independent of assumptions about matter's nature.
• Microscopic Approach (Statistical Thermodynamics):
  • Views matter as composed of a large number of atoms and molecules.
  • Studies thermodynamics requiring knowledge on individual particle behavior.
  • Deals with the effects of molecule action, imperceptible to human senses.
  • Observations depend on the assumptions regarding the nature of matter.
  • Examples include individual molecules present in air.

Continuum

• Continuum disregards the atomic structure of substances, perceiving them as continuous and homogenous.
• Properties are treated as point functions that vary continually.
• Acceptable as long as the system's size is much larger than the space between molecules.
• Exemplified by statements like water density being the same across a glass of water.

Thermodynamic System

• A system is the selected area or matter being studied.
• Surroundings encompass the mass or region beyond the system.
• The boundary demarcates the system from its surroundings
• Systems are classified into closed, open, and isolated systems.

Closed System

• The closed system, also a "control mass," encloses a fixed mass amount which cannot cross its border.
• Mass neither enters nor exits a closed system
• Energy, in the form of heat or work, can pass through the boundary
• The volume of a closed system does not have to be fixed

Open System

• Both mass and energy may cross the boundary of a control volume/open system.
• control volume

Isolated System

• Neither energy nor mass can cross the border.

Properties of a System

• Characteristic of system referenced to as property (pressure P, temperature T, volume V, and mass m).
• Properties can be intensive or extensive
• Intensive properties are independent of the mass of a system (temperature, pressure, and density).
• Extensive properties depend on the size or extent of the system.
• Examples of extensive properties are Total mass, total volume, and total momentum.
• Dividing the system into two equal parts will yield the same value of property (intensive) or half the value (extensive).

State, Path, Process and Cycle

• State A set of properties that completely describes the condition Process Any change that a system undergoes from one equilibrium state to another.
• Path series of states through which a system passes during a process
• Cycle When a system returns to its initial state at the end a process.

Thermodynamic Equilibrium

• Thermodynamics deals with equilibrium states or a state of balance.
• In an equilibrium state, there are no unbalanced potentials within the system.
• The system experiences no changes if isolated.
• Types of equilibrium include:
  • Thermal Equilibrium: Temperature is the same throughout the system.
  • Mechanical Equilibrium: No unbalanced forces thus pressure is constant.
  • Chemical equilibrium : no transfer of mass or chemcial reaction.

Quasi-Static / Quasi-Equilibrium Process

• The system remains close to an equilibrium state.
• The process is sufficiently slow.
• Idealized, not an actual process.
• Approximates actual processes; can be modeled as quasi-equilibrium.
• If weight is taken off of a system slowly, the change is infinitesimally small. This makes it an equilibrium state.
• Joined by all equilibrium states.

Path & Point Functions

• Path functions depend on the path during a process as well as the end states (heat and work).
• Point functions depend on state only (properties).

Different Forms of Energy

• Potential energy (PE): results from its elevation in gravitational field - PE = mgz.
  • g = gravitational acceleration
  • z = elevation of the center of gravity of a system
• Kinetic energy (KE): as a result of its motion relative to some reference frame
  • KE = m V2/2 (when all parts of a system move with the same velocity)
• Microscopic energy forms: relate to the molecular structure of system (independent of outside reference frames).
• Internal energy of system is the sum of the microscopic forms of energy, denoted by U
• Total energy: E = U + KE + PE = U+m V2/2 + mgz

Heat

• Heat transfers energy between two systems (or a system and its surroundings).
• Depends on temperature difference.
• No heat transfer when two systems are at the same temperature.
• Specific Heat. Energy to raise the temperature of a unit mass of a substance by one degree
• Two kinds of specific heats:
  • constant volume Cv
  • constant pressure Cp.
• The specific Heat multiplied by its mass is called Specific Heat capacity

Work

• Energy that crosses a closed system
• Work is the power per unit time denoted by W.
  • Measured in J/s or W.
• Work completed by a system is positive, on a system is negative.
• Heat and Work similarities:
  • recognized at the system boundaries, are boundary phenomena.
  • recognized at the boundaries of a system
  • cross the boundaries
• In general:
  • Systems possess energy, but not heat or work.
  • Associated with a process, not a state.
  • path functions.
• Work is calculated using the formula δW = P dV

Types of Work

• Work is interpreted to be work done at state 1 and state 2, and represented by the area under a curve.
• Volume decreased, meaning the work was done on the system.
• Work is called a path function due to mathematical parlance.

Free Expansion Work

• Work moves at the system boundary
• Free expansion work occurs against a vacuum and has no work transfer.

Zeroth Law of Thermodynamics

• When two bodies have the same temperature with a third body, they have the same temperature with each other.
• Basis for temperature measurement.

Concept of Heat

• Transfer of heat into a system = heat addition
• Transfer of heat out of a system = heat rejection.
• An adiabatic process contains no heat transfer.

Heat Transfer

• Heat transfer depends on the path the system follows as well as on end states
• Inexact differential.
• A graph diagrams that give transfers in two processes.

Temperature Scales

• Temperature scales enable a same basis for measurements.
• Reproducible states like the freezing and boiling points of water:
  • "ice point"
  • "steam point"
• A thermodynamic temperature scales independently of substance properties
  • Kelvin.
• The thermodynamic temperature scale in the English system Rankine.
• Celisus, Kelvin, Rankine:
• T(K) = T(°C) + 273.15
• T(R) = T(°F) + 459.67
• T(R) = 1.8T(K)
• T(°F) = 1.8T(°C) + 32
• Value chosen chose for original Kelvin scale
  • 273.15k (or 0C) water freezes.

Measurement of Temperature Thermometry

• 0th law is called thermometer Thermal equilibrium of bodies calibrated at standard temp. 'body 1 brought in thermal communication
  • 1 has equality with thermometer, and body 2
  • Body "1" has temperature of body "2" example by measure of mercury volume.
• Height of mercury = thermometric property.
  • constant Volumes gas Pressure (p)
  • Constant pressure gas Volume (V)
  • Alcohol or mercury-in-glass Length (L)
  • Electric resistance Resistance (R)
  • Thermocouple Electromotive Force (F)
  • Radiation (pyrometer) Intensity of radiation (I or J)