Physical%20Transformations%20of%20Pure%20Substances

1
Physical Transformations of Pure Substances

• Chapter 4

2
Stabilities of Phase

• A phase of a substance is a form of matter that
  is uniform throughout in chemical composition and
  physical state.
• A phase transition is the spontaneous conversion
  of one phase into another.
• Phase transitions occur at a characteristic
  temperature and pressure.

3
Stabilities of Phase

• At 1 atm, lt 0 C, ice is the stable phase of H2O,
  but gt 0 C, liquid water is the stable phase.
• The transition temperature, Ttrs, is the
  temperature at which two phases are in
  equilibrium.
• So what happens to Gibbs energy?

4
Stabilities of Phase

• At 1 atm, lt 0 C, ice is the stable phase of H2O,
  but gt 0 C, liquid water is the stable phase.
• The transition temperature, Ttrs, is the
  temperature at which two phases are in
  equilibrium.
• So what happens to Gibbs energy?
• lt 0 C Gibbs energy decreases as liquid ? solid.
• gt 0 C Gibbs energy decreases as solid ? liquid.

5
Stabilities of Phase

• Thermodynamics does not provide information
  regarding the rate of phase change.
• Diamond ? graphite
• Thermodynamically unstable phases that persist
  due to slow kinetics are called metastable
  phases.

6
Phase Diagrams

• Phase boundaries show the values of p and T at
  which two phases coexist in equilibrium.

7
Vapor Pressure

• The pressure of a vapor in equilibrium with a
  liquid is called the vapor pressure.
• The pressure of a vapor in equilibrium with a
  solid is called the sublimation vapor pressure.

8
Boiling Point

• Liquid can vaporize from a liquid surface below
  its boiling point as we learnt from the
  Drinking Bird.
• In an open vessel, the temperature at which the
  vapor pressure equals the external pressure, is
  called the boiling temperature.
• At 1 atm, its called the normal boiling
  temperature, Tb.
• At 1 bar, its called the standard boiling point.
• Normal point of H2O is 100.0 C its standard
  boiling point is 99.6 C.

9
Critical Point

• In a closed rigid vessel, boiling does not occur.
• As the temperature is raised the density of vapor
  increases and the density of the liquid
  decreases.
• When the density of the vapor and liquid phases
  are equal the surface between the two phases
  disappears.
• The temperature at which this occurs is called
  the critical temperature, Tc.
• The vapor pressure at the critical temperature is
  called the critical pressure, pc.

10
Critical Point
11
Melting and Freezing

• The temperature at which, under a specified
  pressure, the liquid and solid phases of a
  substance coexist in equilibrium is called them
  melting temperature.
• The freezing temperature is the same as the
  melting point.
• At 1 atm, the freezing temperature is called the
  normal freezing point, Tf.
• At 1 bar, its called the standard freezing
  point.
• The difference is negligible in most cases.
• The normal freezing point is also called the
  normal melting point.

12
Triple Point

• There is a set of conditions under which three
  different phases of a substance (typically solid,
  liquid and vapor) all simultaneously coexist in
  equilibrium.
• This point is called the triple point.
• For any pure substance the triple point occurs
  only at single definite pressure and temperature.
• The triple point of water lies at 273.16 K and
  611 Pa.

13
Triple Point

• The triple point marks the lowest pressure at
  which a liquid phase can exist.

14
Carbon Dioxide
15
Water
16
Helium
17
Thermodynamics of Phase Transitions

• The molar Gibbs energy, Gm, is also called
  chemical potential, m. Phase transitions will be
  investigated primarily considering the change in
  m.
• Thermodynamic definition of equilibrium At
  equilibrium the chemical potential of a substance
  is the same throughout the sample, regardless of
  how many phases are present.

18
Thermodynamics of Phase Transitions
19
Thermodynamics of Phase Transitions

• At low temperatures, and provided the pressure is
  not too low, the solid phase of a substance has
  the lowest chemical potential and is therefore
  the most stable.
• Chemical potentials change with temperature this
  explains why different phases exist.

20
Temperature Dependence of Phase Transitions

• As temperature increases, chemical potential
  decreases.

21
Melting and Applied Pressure

• Molar volume of solid is smaller than that of the
  liquid.

22
Melting and Applied Pressure

• Molar volume of solid is greater than that of the
  liquid.

23
Melting and Applied Pressure
24
Melting and Applied Pressure

• Calculate the effect on the chemical potentials
  of ice and water of increasing pressure from 1.00
  to 2.00 bar at 0 C. The density of ice is 0.917
  g cm-3 and that of liquid water is 0.999 g cm-3.

25
Melting and Applied Pressure

• Calculate the effect on the chemical potentials
  of ice and water of increasing pressure from 1.00
  to 2.00 bar at 0 C. The density of ice is 0.917
  g cm-3 and that of liquid water is 0.999 g cm-3.

26
Vapor Pressure and Applied Pressure

• When pressure is applied to a condensed phase,
  its vapor pressure rises.
• This is interpreted as molecules get squeezed out
  of the condensed phase and escape as a gas.

27
Vapor Pressure and Applied Pressure
28
Location of Phase Boundaries

• Locations of phase boundaries pressures and
  temperatures - can be located precisely by making
  use of the fact that at when two phases are in
  equilibrium, their chemical potentials must be
  equal

29
Location of Phase Boundaries
30
Location of Phase Boundaries
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Solid-liquid boundary
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Solid-liquid boundary
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Solid-liquid boundary
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Solid-liquid boundary
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Liquid-vapor boundary
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Liquid-vapor boundary
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Liquid-vapor boundary
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Liquid-vapor boundary
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Liquid-vapor boundary
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Solid-gas boundary