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    in the phase diagram for water, indicate the direction that the solid–liquid and liquid–gas coexistence lines will move along the temperature axis after the addition of solute.

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    10.4 Phase Diagrams – Chemistry

    10.4 PHASE DIAGRAMS

    Learning Objectives

    By the end of this section, you will be able to:

    Explain the construction and use of a typical phase diagram

    Use phase diagrams to identify stable phases at given temperatures and pressures, and to describe phase transitions resulting from changes in these properties

    Describe the supercritical fluid phase of matter

    In the previous module, the variation of a liquid’s equilibrium vapor pressure with temperature was described. Considering the definition of boiling point, plots of vapor pressure versus temperature represent how the boiling point of the liquid varies with pressure. Also described was the use of heating and cooling curves to determine a substance’s melting (or freezing) point. Making such measurements over a wide range of pressures yields data that may be presented graphically as a phase diagram. A phase diagram combines plots of pressure versus temperature for the liquid-gas, solid-liquid, and solid-gas phase-transition equilibria of a substance. These diagrams indicate the physical states that exist under specific conditions of pressure and temperature, and also provide the pressure dependence of the phase-transition temperatures (melting points, sublimation points, boiling points). A typical phase diagram for a pure substance is shown in Figure 1.

    Figure 1. The physical state of a substance and its phase-transition temperatures are represented graphically in a phase diagram.

    To illustrate the utility of these plots, consider the phase diagram for water shown in Figure 2.

    Figure 2. The pressure and temperature axes on this phase diagram of water are not drawn to constant scale in order to illustrate several important properties.

    We can use the phase diagram to identify the physical state of a sample of water under specified conditions of pressure and temperature. For example, a pressure of 50 kPa and a temperature of −10 °C correspond to the region of the diagram labeled “ice.” Under these conditions, water exists only as a solid (ice). A pressure of 50 kPa and a temperature of 50 °C correspond to the “water” region—here, water exists only as a liquid. At 25 kPa and 200 °C, water exists only in the gaseous state. Note that on the H2O phase diagram, the pressure and temperature axes are not drawn to a constant scale in order to permit the illustration of several important features as described here.

    The curve BC in Figure 2 is the plot of vapor pressure versus temperature as described in the previous module of this chapter. This “liquid-vapor” curve separates the liquid and gaseous regions of the phase diagram and provides the boiling point for water at any pressure. For example, at 1 atm, the boiling point is 100 °C. Notice that the liquid-vapor curve terminates at a temperature of 374 °C and a pressure of 218 atm, indicating that water cannot exist as a liquid above this temperature, regardless of the pressure. The physical properties of water under these conditions are intermediate between those of its liquid and gaseous phases. This unique state of matter is called a supercritical fluid, a topic that will be described in the next section of this module.

    The solid-vapor curve, labeled AB in Figure 2, indicates the temperatures and pressures at which ice and water vapor are in equilibrium. These temperature-pressure data pairs correspond to the sublimation, or deposition, points for water. If we could zoom in on the solid-gas line in Figure 2, we would see that ice has a vapor pressure of about 0.20 kPa at −10 °C. Thus, if we place a frozen sample in a vacuum with a pressure less than 0.20 kPa, ice will sublime. This is the basis for the “freeze-drying” process often used to preserve foods, such as the ice cream shown in Figure 3.

    Figure 3. Freeze-dried foods, like this ice cream, are dehydrated by sublimation at pressures below the triple point for water. (credit: ʺlwaoʺ/Flickr)

    The solid-liquid curve labeled BD shows the temperatures and pressures at which ice and liquid water are in equilibrium, representing the melting/freezing points for water. Note that this curve exhibits a slight negative slope (greatly exaggerated for clarity), indicating that the melting point for water decreases slightly as pressure increases. Water is an unusual substance in this regard, as most substances exhibit an increase in melting point with increasing pressure. This behavior is partly responsible for the movement of glaciers, like the one shown in Figure 4. The bottom of a glacier experiences an immense pressure due to its weight that can melt some of the ice, forming a layer of liquid water on which the glacier may more easily slide.

    Figure 4. The immense pressures beneath glaciers result in partial melting to produce a layer of water that provides lubrication to assist glacial movement. This satellite photograph shows the advancing edge of the Perito Moreno glacier in Argentina. (credit: NASA)

    Source : opentextbc.ca

    Chemistry 1 Exam Flashcards

    Chapters 10 and 11 Learn with flashcards, games, and more — for free.

    Chemistry 1 Exam

    Describe the relative densities of the phases for most substances

    density of gas phase __ density of liquid phase__density of solid phase

    Complete the table describing the shape and volume of each shape

    Click card to see definition 👆

    <

    ~~

    Click again to see term 👆

    Arrange the molecules by strength of the London (dispersion) force interactions between molecules.

    Click card to see definition 👆

    Strongest London dispersion forces

    -CH3 CH2 CH2 CH2 CH2 CH2 CH3

    -CH3 CH2 CH2 CH2 CH3

    -CH3 C(CH3)2 CH3

    Weakest London dispersion forces

    Click again to see term 👆

    1/93 Created by madisonjasper21PLUS Chapters 10 and 11

    Terms in this set (93)

    Describe the relative densities of the phases for most substances

    density of gas phase __ density of liquid phase__density of solid phase

    Complete the table describing the shape and volume of each shape

    <

    ~~

    Arrange the molecules by strength of the London (dispersion) force interactions between molecules.

    Strongest London dispersion forces

    -CH3 CH2 CH2 CH2 CH2 CH2 CH3

    -CH3 CH2 CH2 CH2 CH3

    -CH3 C(CH3)2 CH3

    Weakest London dispersion forces

    Arrange the compounds by boiling point

    pentane: H3C-CH2-CH2-CH2-CH3

    CH3

    neopentane: H3C-C-CH3

    CH3

    hexane: H3C-CH2-CH2-CH2-CH2-CH3

    Highest boiling point

    -hexane -pentane -neopentane

    Lowest boiling point

    Which of the following substances have polar interactions (dipole-dipole forces) between molecules?

    -ClF -NF3

    Which substances exhibit only London (dispersion) forces?

    -Cl2 -He

    Which molecules can hydrogen bond?

    -HF -CH3OH

    Classify each substance based on the intermolecular forces present in that substance

    Hydrogen bonding, dipole-dipole, and dispersion

    -HF

    Dipole-dipole and dispersion only

    -HCl -CO

    Dispersion only

    -CO2

    In the context of small molecules with similar molar masses, arrange the intermolecular forces by strength

    Strongest

    -hydrogen bonding

    -dipole-dipole interactions

    -London dispersion forces

    Weakest

    Arrange these compounds by their expected boiling point

    Highest boiling point

    -CH3OH -CH3Cl -CH4

    Lowest boiling point

    What would happen to each of the properties if the intermolecular forces between molecules increased for a given fluid? Assume temperature remains constant

    Increase

    -boiling point -viscosity -surface tension

    Decrease

    -vapor pressure

    For liquids, which of the factors affect vapor pressure?

    -intermolecular forces

    -temperature

    Arrange these compounds by their expected vapor pressure

    Highest vapor pressure

    -Br2 -NCl3 -H2O

    Lowest vapor pressure

    The heat of vaporization of water is 40.66 kJ/mol. How much heat is absorbed when 3.00 g of water boils at atmospheric pressure?

    heat: 6.769 kJ

    Classify each phase change based on whether it describes a transition between a gas and a liquid, a gas and a solid, or a liquid and a solid

    Gas and liquid

    -condensation -evaporation

    Gas and solid

    -sublimation -deposition

    Liquid and solid

    -freezing -melting

    Label the heating curve with the phase or phases present. Assume constant pressure

    Calculate the energy released when 19.1 g of liquid mercury at 25.00*C is converted to solid mercury at its melting point

    Constants for mercury at 1 atm

    heat capacity of Hg(l) :28.0J/(mol x K)

    melting point: 234.32K

    enthalpy of fusion: 2.29 kJ/mol

    0.387 kJ

    At 1 atm, how much energy is required to heat 43.0 g H2O(s) at -20.0C to H2O(g) at 141.0C? Use the heat transfer constants found in this table.

    134.7 kJ

    Two 20.0 g ice at -12.0°C are placed into 265 g of water at 25*C. Assuming no energy is transferred to or from the surroundings, calculate the final temperature, Tf, of the water after all the ice melts

    heat capacity of H2O(s): 37.7J/(mol x K)

    heat capacity of H2O(l): 75.3J/(mol x K)

    enthalpy of fusion of H2O: 6.01kJ/mol

    ????

    In what phase is CO2 at 25 atm and -65*C?

    Source : quizlet.com

    Answered: In the phase diagram for water,…

    Solution for In the phase diagram for water, indicate the direction that the solid–liquid and liquid-gas coexistence lines will move along the temperature axis…

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    Transcribed Image Text:In the phase diagram for water, indicate the direction that the solid–liquid and liquid-gas coexistence lines will move along the temperature axis after the addition of solute. Answer Bank liquid solid gas Temperature (°C) Pressure (atm)

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