Chapter 2 Physical Properties of Pure Componds

Each boiling point corresponds to a vapor pressure

where saturated liquid and vapor can exist in equilibrium with each other

liquid and solid can exist in equilibrium with each other at any melting or freezing point

Saturated vapor

a vapor at a boiling point

saturated liquid

a liquid at a boiling point

Superheated vapor

a vapor at a T above its boiling point

Subcooled liquid

a liquid at a temperature below its boiling point

compressed liquid

A liquid at a pressure above its vapor pressure

State properties

a property that describes the condition of a material or system at a particular time

the value of the state property is independent of how the system arrived (path taken) AT ITS CURRENT STATE

p,v,t,n

heat and work are not- path dependent

Intensive properties

do not depend on wither the size of the system or the amount of material

density or viscosity

Extensive properties

are proportional to the amount of material

volume, mass, enthalpy

Gibbs Phase rule

F= C- pi + 2

Degrees of freedom are the numbers of intensive variable that can be established before all others are constrained to unique values

F

number of degrees of freedom

Pi

number of distinct phases

C

number of distinct chemcial compounds

Quality (q)

of a vapor liquid system us the mass or mole fraction of the vapor phase

q= mass vapor/ mass system

If the liquid and vapor phases are each made up of the same pure compound then M=N

If a system consists of liquid and vapor in equilibrium

an intensive property of the overall system can be computes as

X=(1-q)X^L +qX^v

X is any intensive property

Enthalpy

H= U+pv

pv- flow work J

state property that quantifies energy

quantified to a reference state

extensive property, but has intensive counterparts like molar enthalpy and specific enthalpy

Heat capacity

an intensive property

constant volume heat capacity (Cv)

and constant pressure heat capacity (Cp)
look over equations

dU= Cv dT

is valid when molar volume is constant

Always valid for ideal gases

dH = Cp dT

if pressure is constant

always valid for ideal gases

Ideal gas law

is a hypothetical gas in which the molecules have no intermolecular interactions and no volume

Works better for lower pressure systems

Both molar internal energy and enthalpy of an ideal gas

are only dependent on temperature

Higher temp has more energy

For ideal gases

Cv + R = Cp

Monatomic ideal gas

Cv= (1.5)R

Equation of state

a relationship among T, p, and V such that if any two are known the third can be calculated

ex. the ideal gas law

van der waals

ideal gas behavior

p=RT/V-b - a/V²

a and b are constants that have unique values for each compound

a- effect of intermolecular attractions

b- volume of the particles

van der waals

If V >> b, then (V – b) ~ V
• If V is very large, the a/V2 term will be negligible.
• For a very large V, van der Waals EOS reduces to the ideal gas law
• Physically, when V is large, pressure is correspondingly low: well
modeled as ideal gases

Compressibility factor Z

Z=PV/RT

Z = 1 for an ideal gas
• EOS are frequently expressed in terms of Z to make comparisons
to ideal gas behavior easier
• The van der Waals EOS written in terms of Z:

z= V/V-b - a/RTV

For most liquids and solids

molar volume (or specific volume) changes little with changes in temperature and pressure

Both liquids and solids are frequently modeled as having constant volume
(especially over small ranges of T and/or P)

𝑉 is much smaller in solids and liquids than in vapors and gases. At low
pressure, P𝑉 is usually very small compared to 𝑈.

For solids and liquids at low pressure

𝐻 = 𝑈 + 𝑃𝑉 ≈ 𝑈 , 𝐶𝑉 ≈ 𝐶𝑃