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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.
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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
