CME 304 Definitions

Quiz 9/8

Thermodynamics

the study of energy transformations associated with material substances and of how these bodies are affected by the transport of different forms of energy into and out of them

Matter

• all matter exists in the form of phases

  • solids, liquids, vapors, plasmas, supercritical fluids

• a pure phase (atom/molecule) has a distinct molecular arrangement and chemistry

  • is homogeneous and isotropic throughout

  • is separated from other phases by a distinct boundary surface

    • cannot be a mizture of two or more phases of differing composition

    • exception = critical state of matter

• phases of matter may exist in different forms

  • pure carbon exists as two solid phases (graphite/diamond)

  • He (two liquid phases, ³He and ⁴He)

  • Fe (three solid phases, one liquid phase, one vapor phase)

  • H₂O (seven solid phases, one liquid phase, one vapor phase)

Solids

• falls under “Intermolecular Forces in Pure Phases”

  • atoms or molecules arranged in a 3-dimensional network that is repeated in a fixed pattern (exception: glass)

  • attractive forces between atoms or molecules are strong enough to keep them in fixed equilibrium positions where there is a balance of attractive and repulsive forces

  • atoms or molecules oscillate (or vibrate) about their fixed equilibirum positions

  • magnitude of oscillations depends on temperature

  • when oscillations become too great, intermolecular forces are not strong enough to hold structure together and the —— tears apart and becomes a liquid

    tldr: increased temperature increases oscillations, tearing apart the structure and changing its phase

liquids

• falls under “Intermolecular Forces in Pure Phases”

  • intermolecular spacings are of the same order as in solids

  • atoms or molecules are no longer in fixed positions relative to each other

  • atoms or molecules can rotate and translate and vibrate more freely than in case of solids

  • intermolecular binding forces are weaker than in solids but still strong compared to gases

  • molecular volume is generally larger than solids

    • (exception: liquid water and ice)

tldr: less strong IMF but still some structure, can rotate, translate, and vibrate more than solids but less than gas

gases

• falls under “Intermolecular Forces in Pure Phases”

  • molecules are far apart leading to lower densities and larger molecular volumes

  • particles move in random fashion

  • particles continually collide with each other and with the container walls

  • weak IMF (except at high pressures)

  • particles have high kinetic energy level compared to liquids and solids

  • molecules can rotate and vibrate within

  • must release a large amount of energy to condense into liquid or solid phases

tldr: molecules are far apart and free to move, continuous collisions with each other and container walls = increased/high kinetic energy

Energy (in thermodynamics)

• a characteristic (or property) of a finite material body in equilibrium with its surroundings that gives it the capacity to convey some portion of this characteristic via thermal (heat) and/or mechanical (work) means to its immediate surroundings (or vice versa)

• thermodynamics is the study of how the energy of such a material body can interact in prescribed ways with its surroundings (i.e., across a physical boundary) to transfer work and/or heat

• a body can change its physical or chemical state due to such an interaction with its surroundings

  • heating of liquid water (change of temperature/pressure), melting of ice (change of phase), expansion/compression of a fluid (change of shape), combustion of chemical fuel (change of chemical species), dissolution of sugar in water (change of chemical composition)

Energy Storage

• a fundamental principle is that a material body can store its energy in different forms

  • electronic

  • nuclear

  • kinetic

  • vibrational

  • rotational

  • magnetic

  • gravitational

  • surface

• it is assumed that any interaction of a material body with its surroundings could result in the changes in the relative amounts of energy storage

Energy Transformations

• fundamental principle of thermodynamics

  • Different forms of energy can be transformed or exchanged with each other

  • i.e.

    • electrical to thermal (light bulb)

    • mechanical to electric (power station)

    • chemical to thermal (combustion of a fuel)

    • thermal to mechanical (auto piston)

    • nuclear to thermal (nuclear reactor)

Property

• coordinate

  • a macroscopic characteristic of a finite body of matter that describes the internal state of the system

  • a specific numerical value can be assigned without knowledge or reference to its previous history

    • it is time independent

    • uniform throughout (homogeneous)

    • independent of direction of measurement (isotropic)

    • independent of path chosen

    • typical properties might be:

      • pressure

      • volume

      • temperature

tldr: often the thing we are solving for or using to solve for something else, it describes the internal state of the system and can be a numerical value

Absolute Property

• does NOT depend on a choice of reference state

  • Density

  • Heat Capacity

  • Thermal Expansion

  • Compressibility

  • Electrical conductivity

Floating property

• properties that cannot be measured directly but must be computed relative to some arbitrarily assigned reference condition

  • there are no meters or devices for measuring ———- ————

  • only property changes can be determined

  • property value depends on choice of reference state

  • if we arbitrarily set U (internal energy) = 0 at T = 0 degrees C and P = 1 bar, then all calaculations of internal energy are made relative to this reference condition

  • since we are generally more interested in determining property changes than in absolute property values, the choice of reference condition does not matter since it cancels out

tldr: the property we calculate the change of, not the actual value, need reference conditions in order to do this

examples:

• internal energy

• enthalpy

• entropy

• Gidds Free Energy

• Gravitational potential energy

• Chemical potential

• Kinetic energy

*most thermodynamic properties are ——— ——— in they “float” depending on choice of reference state

Extensive property

• DEPEND on the SIZE of the system

  • if a system is sub-divided into smaller units and a system property is identified then the extensive value of that property is the sum of all the sub-units that make up the system (i.e. volume, area, mass)

  • extensive properties can change with time as a system interacts with its surroundings

tldr: key word DEPEND

Intensive property

• INDEPENDENT of the size of the system and may vary from place to place within the system at any moment of time

  • can vary with both position and time as a system interacts with its surroundings

  • ex: pressure, temeprature, heat capacity, density, specific or molar volume, volume expansion coefficient, compressibility, chemical composition, specific internal energy or enthalpy

tldr: key word INDEPENDENT

Non-property

• a quantity (or characteristic) is a ——— ———- if its change in value between two states depends on the details of the process and not solely on the end states themselves

  • heat transfer between two bodies at different temperatures

  • mechanical energy transfer between two different bodies (work)

tldr: its value depends on the process more than the end states of it, unlike a property, it is not independent of the details of the process

Thermodynamic state of a body

• the sum totality of all its properties

• since there are often known mathematical relations between different properties, the thermodynamic state can often be described as a subset of properties from which other properties can be determined

• thermodymic properties are uniform throughout body

• thermodynamic properties are time independent

control properties

• properties which can be externally prescribed to alter the energy state of a materials system

  • pressure

  • temperature

  • volume

  • chemical composition

dependent properties

• properties that are determined (or dependent on) the setting of controllable properties

  • internal energy, enthalpy, entropy, free energy

  • heat capacity (at cont T or V)

  • volume thermal expansion

  • compressibility

macrostate

• any state where the thermodynamic properties can be easily measured using laboratory equipment (i.e. P,T, V)

• we will be conerned here only with the thermodynamic behavior macroscopic systems since the study of microscopic is beyond the current scope of this discussion

microstate

• any state where the thermodynamic properties are determined by microscopic quantum mechanical parameters

• this is the realm of statistical thermodynamics and is beyond the scope of this course

equtions of state

• it has been shown that a mathematical relationship exists among the thermodynamic properties (P,V,T) that specify the equilibrium thermodynamic state of the system

  • for an ideal gas we need 3 properties (i.e. P,V,T)

    • PV = nRT

    • VanderWaals gas

      • (P+a/V²)(V-b)=NkT

  • thus if we specify any two thermodynamic properties, the ———- — —— will, in principle, determine the third

  • thus the thermodynamic state of a system can always be specified by any two thermodynamic properties and the ——— — —— yields the 3^rd

tldr: essentially the principle that if you are given two properties, this can be used to determine the third, think physics equations

thermodynamic system

• a system is any object, or finite quantity of matter that occupies a region of space that is selected to be set aside for study

• the system is treated as a whole unit

  • the system can be described by a set of properties that apply to the system as a whole

• system properties can be changed by interaction with surroundings

  • system may exchange energy or material flow with surroundings

  • the # of system properties used to describe the system is small (i.e. 3-6)

*make sure you KNOW THIS

sub-system

• a ———- is obtained when a thermodynamic system is divided into two or more sub-parts separated by a sub-system boundary and each subsystem treated as a separate entity as shown below

• any property changes in the system is the result of summing the property changed in each sub-system that make up the system

macroscopic system

• a system that is treated as a whole unit

• the paramters that describe the system must apply to system as a whole (i.e. P, T, V, X=composition)

• the number of thermodynamic coordinates used to describe system is small (i.e. 3-6)

microscopic system

• a system treated as a collection of minute atomic/molecular discrete entities

• system varibles apply only to individual particles

• number of coordinates need to describe system is very large (10²³)

• relation of microscopic to macroscopic system properties is treated in statistical thermodynamics

surroundings

• the portion that is NOT the system

• only includes that portion of the immediate space in the neighborhood of the system that is capable of interacting with the system

  • can absorb heat from or release heat to the system through boundary

  • can have work performed on it by the system or can do work on the system

  • can absord matter from the sysbtem or release matter into the system

*KNOW THIS

system boundary

• we assume that a distinct physical and/or chemical ——— (or interface) forms at the periphery of the system and separates the system from its surroundings

• may or may not allow the system to interact with its surroundings

  • if it permits heat transfer and mechanical energy transfer into or out of the system but NOT mass flow, it is CLOSED

  • if it does NOT permit heat transfer into or out of the system, it is an adiabatic boundary

  • if it permits heat transfer, mechanical energy transfer and mass flow into or out of the system it is OPEN

  • if it is fixed then it is a rigid ———- (rigid container): otherwise if the ———- can be displaced it is a moveable ———— (piston; cylinder)

  • if an open system has particular locations where it can interact with its surroundings, it is either an entrance or exit port ocated at the —————

*KNOW THIS

universe

• includes both the system + surroundings

• _________ = system + surroundings

• any changes (i.e. energy, entropy) in the _______ must be the result of the combined changes in the System and Surroundings

• thus, we may write:

• delta(______) = delta(System) + delta(Surroundings)

*KNOW THIS

types of thermodynamic systems

• there are three types of systems that we will consider:

  • closed system

  • isolated system

  • open system (control volume)

• see further cards for definitions

closed system

• system which is in thermal or mechanical contact with its surroundings but does NOT allow the passage of matter into or out of it

  • heat and/or work can be performed on the system by surroundings or vice versa

isolated system

• system is NOT in thermal, mechanical, or physical contact with the surroundings

  • e.g. a thermos bottle

open system

• aka control volume system

• system not only is in thermal and mechanical contact with its surroundings but also will allow the passage of matter through it

*think of a nozzle

thermodynamic equilibrium

• thermodynamics is concerned with equilibirum states of matter

• implies a state of matter in which all the forces acting on a system are in balance

• there are four basic types of equilibrium that are of interest

  • mechanial equilibrium

  • thermal equilibrium

  • phase equilibrium

  • chemical equilibrium

• criteria:

  • isolate system from surroundings

  • look for measureable or observable changes in system properties

  • if no changes occur over time, the system is in eq at the time of isolation and is said to be in its equilibrium state

• there is NO requirement that system be in an equilibrium state during changes of state

mechanical (hydrostatic) equilibrium

• no unbalanced forces acting on the system

• system is stationary in time and space

thermal equilibrium

• no temperature gradients throughout system

• system must be at uniform temperature

• no net heat flow between different systems

chemical equilibrium

• no chemical reactions take place at measurable rates

• rate of forward reactions = rate of reverse reactions

phase equilibrium

• two or more different phases of matter (solid, liquid, vapor) co-exist at a given P and T

driving forces for equilibrium

• changes of state occur bc of the presence of an imbalance in driving forces that cause the system to move from one equilibrium state to another

• these driving forces (or constraints on the system) are the result of:

  • temperature differences leading to flow of heat between systems and surroundings

  • pressure differences leading to expansion or contraction of system/surroundings

  • compositional differences leading to the diffusion chemical species between regions of differing composition

  • chemical reactions between different reacting species to produce new chemical species

  • temperature, pressure and/or compositional difference that lead to the formation of a new phase

  • exception: the presence of energy barriers can prevent equilibrium from occurring in some systems

• we conclude that the system is at equilibrium when all driving forces are uniform throughout system and surroundings and that all barriers to change have been overcome

equilibrium postulates

• all thermodynamic systems will spontaneously tend to move toward their equilibrium states

  • the equilibrium state will be determined by the constraints imposed on the system (P,T,V,X) where X=composition

  • the presence of thermodynamic energy barriers may prevent equilibrium from being attained

    • leads to metastable equilibrium states

• Classical Thermodynamics is based upon the determination of the equilibrium states of matter

  • all changes of state are assumed to be from one equilibrium state to another equilibrium state

  • once the equilibrium state has been achieved, the prior history of the system is of no relevance

thermodynamic process

• takes place when a system is forced to leave its equilibrium state due to interaction with the surroundings and move toward another equilibrium state dictated by the new set of constraints imposed on it

• any such changes of equilibrium condition is referred to as a change of state

• two different states must be distinguishable by having measurably different properties

• all thermodynamic properties are independent of the choice of path

• thermodynamics is the study of processes leading to a change in thermodynamic properties of a system

path

• a specific pre-selected process whereby a system changes its properties according to a prescribed set of parameters that describe the path

• a system can change its state via different paths

• a system can move along an open path or a closed path

  • an open path does not allow the system or return to its original starting point (state)

  • a closed path allows the system to return to its original starting point

constraint

• a constraint on a thermodynamics process occurs when a thermodynamic property is fixed at a pre-selected value and the thermodynamic process is carried out with this fixed value throughout

• most thermodynamic processes are carried out with at least one constraint:

  • const. temp

  • const. pressure

  • const. volume

  • const. chemical composition

  • no heat flow

  • no mass flow

infinitesimal process

• a process carried out such that only infinitesimal changes in thermodynamic coordinates can occur (i.e. dT, dV, dP)

finite process

• takes place where finite changes in thermodynamic coordinates can be measured (deltaT, deltaP, deltaV)

quasi-static process

• a process carried out so slowly that the system is always close to thermal, mechanical, or chemical equilibrium

• rate of change of process must be much slower than the time it takes for system to respond to external changes (e.g., slow withdrawal of a frictionless piston inside of piston-cylinder assembly containing a gas)

isothermal

• type of thermodynamic process

• process carried out such that the system temperature remains constant throughout

adiabatic

• type of thermodynamic process

• process carried out such that no heat is allowed to enter or escape the system (isolated)

isobaric

• type of thermodynamic process

• process carried out such that the system pressure remains constant throughout

isometric

• type of thermodynamic process

• process carried out at constant volume

reversible process

• a process in which a system and all parts of the surroundings can be restored to their initial states after a given process is reversed along the same path

  • ________ processes are idealized processes that do not occur in nature

  • ________ processes are defined to allow making important thermodynamic calculations and comparisons

• delta(Universe) = delta(System) + delta(Surroundings)=0 or

• delta(System) = -delta(Surroundings)

• can be thought of as the limiting case in which irreversibilities within the system or surroundings are reduced to their lowest values

  • all attempts to reduce spontaneous changes within the system will increase reversibility

irreversible process

• a process in which the system or surroundings. (or both) are permanently changed when the process is reversed from the final to initial state

  • they occur in most naturally occuring changes

  • difficult the quantiy accurately

  • delta(Universe) = delta(System) + delta(Surroundings) =/ 0

• examples

  • unrestrained heat transfer between heat conducting bodies with finite temperature differences

  • unrestrained expanson of a gas at high pressure to a lower pressure

  • spontaneous chemical reaction (combustion)

  • spontaneous mixing of materials of two different compositions

  • friction due to rubbing of surfaces

  • electric current flow through a resistance

  • a chemical explosion

  • cracking of an egg

  • erosion of a mountain by wind and rain

summary of reversible processes

• processes that are carried out in which no barrier to system movement from one equilibrium state to another are present

• system is never more than differentially removed from equilibrium state

• process is quasi-static

• process traverses through a succession of equilibirum states

• driving forces are differential in magnitude

• process can be reversed by reversing direction of differential driving forces

• initial system state is restored after reverse process is completed

internal irreversibilities

• those irreversible processes that occur within the system boundary

external irreversibilities

• those processes that occur outside the system boundary (i.e. surroundings)

equilibrium and reversibility postulate

• all processes that deviate differentially from their equilibrium state are considered to be reversible

  • processes that deviate differentially from their equilibrium state are assumed to be so close to equilibrium that the likelihood of irreversibilities creeping in is thought to be very small

  • if system energy at equilibrium is U₀, then a disturbance in, say, temperature of +dT might bring the system into a new state, U₀+dU, such that the system can be returned infinitesimally and reversibly to its initial state by simply reversing the direction of temperature, -dT