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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

### 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 👆

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~~

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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

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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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