Physical Transformations of Pure Substances

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# Physical Transformations of Pure Substances - PowerPoint PPT Presentation

Physical Transformations of Pure Substances. Chapter 4. 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.

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### Physical Transformations of Pure Substances

Chapter 4

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.
Stabilities of Phase
• At 1 atm, < 0 °C, ice is the stable phase of H2O, but > 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?
Stabilities of Phase
• At 1 atm, < 0 °C, ice is the stable phase of H2O, but > 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?
• < 0 °C Gibbs energy decreases as liquid  solid.
• > 0 °C Gibbs energy decreases as solid  liquid.
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.
Phase Diagrams
• Phase boundaries show the values of p and T at which two phases coexist in equilibrium.
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.
Boiling Point
• Liquid can vaporize from a liquid surface below it’s 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, it’s called the normal boiling temperature, Tb.
• At 1 bar, it’s called the standard boiling point.
• Normal point of H2O is 100.0 °C; it’s standard boiling point is 99.6 °C.
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.
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, it’s called the standard freezing point.
• The difference is negligible in most cases.
• The normal freezing point is also called the normal melting point.
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.
Triple Point
• The triple point marks the lowest pressure at which a liquid phase can exist.
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.
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.
Temperature Dependence of Phase Transitions
• As temperature increases, chemical potential decreases.
Melting and Applied Pressure
• Molar volume of solid is smaller than that of the liquid.
Melting and Applied Pressure
• Molar volume of solid is greater than that of the liquid.
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.
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.
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.
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