chapter 9 liquids and solids · change of state from solid to liquid to gas does not involve change...

Chapter 9

Liquids and Solids

Chapter 9Table of Contents

Copyright ©2016 Cengage Learning. All Rights Reserved.

(9.1) Intermolecular forces

(9.2) The liquid state

(9.3) An introduction to structures and types of solids

(9.4) Structure and bonding in metals

(9.5) Carbon and silicon: Network atomic solids

(9.6) Molecular solids

(9.7) Ionic solids

(9.8) The various forces in substances: Summary and comparison

(9.9) Vapor pressure and changes of state

(9.10) Phase diagrams

Chapter 9


Questions to Consider

Why is white phosphorus stored in water and not left out in the open like sulfur?

What causes the “steaming” of a piece of dry ice?

Section 9.1Intermolecular Forces


An Introduction

Intramolecular bonding - Atoms form molecules by sharing electrons

Bonding within the molecule

Stronger than intermolecular forces

Liquids and solids are the condensed states of matter

They are formed by forces that may involve:

Covalent bonding

Ionic bonding

Intermolecular forces



Figure 9.1 - Schematic Representations of the Three States of Matter



Phase Changes

Change of state from solid to liquid to gas does not involve change in molecules, but change in forces among molecules

Solid to liquid

By adding energy, the motion of molecules increase

Molecules eventually have greater movement and the disorder characteristics of a liquid

Liquid to gas

By adding more energy, individual molecules move far apart, reducing interaction



Table 9.1 - Densities of the Three States of Water



Intermolecular Forces

They are forces that occur between molecules and include:

Dipole–dipole forces

Hydrogen bonding

London dispersion forces



Dipole–Dipole Forces

Exhibited by molecules with polar bonds that behave in an electric field

Attract each other electrostatically

Also called dipole–dipole attraction

Dipoles find the best compromise between attraction and repulsion

Characteristics:

Only 1% as strong as covalent or ionic bonds

Forces grow weaker as distance between dipoles increase



Figure 9.2 - Dipole–Dipole Forces

(a) The electrostatic interaction of two polar molecules

(b) The interaction of many dipoles in a condensed state



Hydrogen Bonding

Strong dipole–dipole forces can be noticed when H is bound to lone pairs on highly electronegative atoms like N, O2, and F

Strength of interactions can be characterized by:

Polarity of the bond

Close approach of the dipoles

Small size of the hydrogen atom

Such bonding affects the physical properties of elements

Important in organic molecules



Figure 9.3 (b) - Hydrogen Bonding Among Water Molecules

Blue dotted lines represent intermolecular forces between the water molecules



Figure 9.4 - The Boiling Points of the Covalent Hydrides of the Elements in Groups 4A, 5A, 6A, and 7A



London Dispersion Forces

They exist among noble gas atoms and nonpolar molecules

Instantaneous dipole that occurs accidentally in a given atom induces a similar dipole in a neighboring atom

Instigates weak and short-lived interatomic attraction

Polarizability - Ease with which the electron cloud of an atom can be distorted to produce dipolar charge distribution

Large atoms with many electrons have high polarizability

Also applies to nonpolar molecules



Figure 9.5 (a) - Polarization



Concept Check

Which of the following is the stronger bond?

a. Intramolecular bond

b. Intermolecular bond

Justify the answer with a valid reason



Concept Check

Identify the molecule that is capable of forming stronger intermolecular forces?

N2 H2O

Justify the choice with a valid reason



Concept Check

Following are the Lewis structures for ethanol and dimethyl ether:

Compare the boiling points of the two molecules

C

H

H C

H

H

H

O H C

H

H C

H

H

H

O H



Concept Check

Which gas would behave more ideally at the same pressure and temperature?

CO N2

Justify the choice with a valid reason

Section 9.2The Liquid State


Characteristics of Liquids

Liquids exhibit low compressibility, lack of rigidity, and high density

Surface tension is the resistance of a liquid to an increase in its surface area

Liquids with large intermolecular forces tend to have high surface tensions

Capillary action is exhibited by polar liquids and involves spontaneous rising of a liquid in a narrow tube

Involves adhesive and cohesive forces



Figure 9.7 - Convex and Concave Meniscus

Nonpolar liquid, mercury, forms a convex meniscus in a glass tube, whereas polar water forms a concave meniscus



Characteristics of Liquids

Viscosity refers to the measure of a liquid’s resistance to flow

Liquids with large intermolecular forces or molecular complexity tend to be highly viscous

Example - Glycerol possess unusually high viscosity because it forms hydrogen bonds using its O–H groups



Structural Model for Liquids

Liquids have strong intermolecular forces and significant molecular motion

Rapid changes in liquids can be studied using spectroscopy

A typical liquid contains a large number of regions with more disorder

The number of regions where holes appear is smaller

Their nature is highly dynamic, rapidly fluctuating in both regions

Section 9.3An Introduction to Structures and Types of Solids


Classification of Solids

Amorphous solids have disordered structures

Example - Glass

Crystalline solids have a highly regular arrangement of components

Positions of components are represented by lattice(s)

A three-dimensional system of points that designates the position of the components of a substance

The smallest repeating unit is termed a unit cell



Figure 9.9 - Three Cubic Unit Cells and the Corresponding Lattices



Figure 9.9 - Three Cubic Unit Cells and the Corresponding Lattices (Contd)



X-Ray Analysis of Solids

X-ray diffraction is used to determine the structure of crystalline solids

Diffraction occurs due to:

Constructive interference when parallel beam waves are in phase

Destructive interference when waves are out of phase

Distance traveled by waves depends on the distance between the atoms

A diffractometer is used to carry out X-ray analysis of crystals



The Bragg Equation

Used in the estimation of interatomic spacing

Where

n is an integer

λ is the wavelength of the X rays

d is the distance between atoms

is the angle of incidence and reflection

λ = 2 sin n d θ



Figure 9.11 - Reflection of X rays of Wavelength λ from a Pair of Atoms in two Different Layers of a Crystal



Interactive Example 9.1 - Using the Bragg Equation

X rays of wavelength 1.54 Å were used to analyze an aluminum crystal. A reflection was produced at = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing this reflection

Solution

To determine the distance between the planes, use the Bragg equation where

n = 1

λ = 1.54 Å

= 19.3 degrees

θ

θ



Interactive Example 9.1 - Using the Bragg Equation

Since 2 sin = λd θ n

o

o1 1.54 A

λ= = = 2.33A = 233 pm

2 sin 2 0.3305

nd

θ



Types of Crystalline Solids

Ionic solids possess ions at the points of the lattice that describe their structures

Molecular solids have discrete covalently bonded molecules at each lattice point(s)

An atomic solid has atoms at the lattice points that describe its structure

Metallic solids

Network solids

Group 8A (18) solids



Figure 9.12 - Examples of Three Types of Crystalline Solids

(a) An atomic solid (b) An ionic solid (c) A molecular solid



Table 9.3 - Classification of Solids

Section 9.4Structure and Bonding in Metals


The Closest Packing Model

Closest packing is an arrangement that assumes that metal atoms are hard, uniform spheres

These spheres are packed in layers

Each successive layer is formed when spheres occupy a dimple formed by the spheres of the previous layer

Forms of closest packing:

aba packing

abc packing



aba Packing

The 2nd layer is like the 1st, but it is displaced, so that each sphere in the 2nd layer occupies a dimple in the 1st layer

The spheres in the 3rd layer occupy dimples in the 2nd layer

Spheres in the 3rd layer lie directly over those in the 1st layer

The resultant structures are called hexagonal closest packed (hcp) structures

The spheres in every layer occupy the same vertical position

When the spheres are aba closest packed, the unit cell is a hexagonal prism



Figure 9.13 (a) - The Closest Packing Arrangement of Uniform Spheres



Figure 9.14 - Hexagonal Closest Packing



abc Packing

The spheres in the 3rd layer occupy dimples in the 2nd layer

Spheres in the 3rd layer do not rest above spheres in the 1st

layer

The 4th layer is like the 1st

The resultant structure is termed a cubic closest packed (ccp) structure

The spheres in every fourth layer occupy the same vertical position

In abc packing, the unit cell is face-centered cubic



Figure 9.13 (b) - The Closest Packing Arrangement of Uniform Spheres



Figure 9.15 - Cubic Closest Packing



Common Characteristics of the hcp and the ccp Structures

Each sphere in both structures possesses 12 equivalent nearest neighbors



Net Number of Spheres in a Face-Centered Cubic Unit Cell - Derivation

A unit cell is defined by the centers of the spheres on the corners of the cube

Net number of spheres in a face-centered cubic unit would be1 1

8× + 6× = 48 2



Interactive Example 9.2 - Calculating the Density of a Closest Packed Solid

Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 144 pm. Calculate the density of solid silver

Solution

Density is mass per unit volume

It is necessary to know how many silver atoms occupy a given volume in the crystal

The structure is cubic closest packed, which means the unit cell is face-centered cubic

Find the volume of this unit cell for silver and the net number of atoms it contains




In the structure, the atoms touch along the diagonals for each face and not along the edges of the cube

The length of the diagonal is r + 2r + r, or 4r

Use this information to find the length along the edge of the cube by the Pythagorean theorem

22 2

2 2

2 2

2

+ = 4

2 = 16

= 8

= 8 = 8

d d r

d r

d r

d r r




Since r = 144 pm for a silver atom,

The volume of the unit cell is d3, which is (407 pm)3, or 6.74× 107 pm3

Converting to cubic centimeters,

= 144 pm 8 = 407 pmd

3 107 3 23 31.00×10 cm

6.74 × 10 pm × = 6.74×10 cmpm




It is known that the net number of atoms in the face-centered cubic unit cell is 4

4 silver atoms are contained in a volume of 6.74 × 10 –23 cm3

Therefore, the density is

23

23 3

3

4 atoms 107.9 g / mol 1 mol / 6.022×10 atomsmassDensity = =

volume 6.74×10 cm

= 10.6 g / cm



Concept Check

Determine the number of metal atoms in a unit cell if the packing is:

a. Simple cubic

b. Cubic closest



Concept Check

A metal crystallizes in a face-centered cubic structure

Determine the relationship between the radius of the metal atom and the length of an edge of the unit cell



Bonding Models for Metals - An Introduction

A successful bonding model for metals must consider:

Malleability

Ductility

Efficient uniform conduction of heat and electricity

Bonding models for metals include:

Electron sea model

Band model (MO model)



Bonding Models for Metals

Electron sea model

A regular array of metal cations are considered to be in a sea of mobile valence electrons

Mobile electrons conduct heat and electricity

Ions can freely move around when the metal is hammered or drawn into a wire



Figure 9.18 - The Electron Sea Model for Metals

(a) Representation of an alkali metal (Group 1A) with one valence electron(b) Representation of an alkaline earth metal (Group 2A) with two valence

electrons



Bonding Models for Metals

Band model, or the molecular orbital (MO) model

Electrons are assumed to be traveling in molecular orbitals formed from the valence atomic orbitals of the metal atom

Interaction between metal atoms results in a virtual continuum of levels, called bands



Figure 9.19 - Molecular Orbital Energy Levels Produced When Various Numbers of Atomic Orbitals Interact



Figure 9.20 - Representation of Energy Levels in a Magnesium Crystal



Metal Alloys

An alloy is a substance that contains a mix of elements and possesses metallic properties

Classification of alloys:

Substitutional alloy, where some of the host metal atoms are replaced by other metal atoms of similar size

Example - Brass

Interstitial alloy, where some of the holes in the closest packed metal structure are occupied by small atoms

Example - Steel



Figure 9.21 - Two Types of Alloys



Carbon and Steel

Amount of carbon affects the properties of steel

Mild steel - Less than 0.2% carbon

Malleable and ductile

Can be drawn into nails, cables, and chains

Medium steel - 0.2 to 0.6% carbon

Harder than mild steel

Used in structural steel beams

High-carbon steel - 0.6 to 1.5% carbon

Used for making cutlery and tools

Section 9.5Carbon and Silicon: Network Atomic Solids


Network Solids - An Introduction

Network solids are those atomic solids that contain directional covalent bonds which form solids that can be viewed as “giant molecules”

Properties:

Brittle nature

Ineffective conductors of heat and electricity

Important elements:

Carbon

Diamond and graphite

Silicon



Forms of Carbon - Diamond

Hardest naturally occurring substance

Each C atom is surrounded by a tetrahedral arrangement of other C atoms

Structure as per the localized electron model:

Stable structure is obtained via covalent bonds

Formed by the overlap of sp3 hybridized C atomic orbitals

Structure as per the molecular orbital theory:

Large gaps between filled and empty levels

Electron transfer is not easy



Figure 9.22 (a) - The Structure of Diamond



Figure 9.23 (a) - Partial Representation of the Molecular Orbital Energies in Diamond



Forms of Carbon - Graphite

Slippery, black, and a conductor of heat and electricity

Structure is based on layers of C atoms arranged in fused six-membered rings

Structure as per the localized electron model:

Shows trigonal planar arrangement

120-degree bond angles

sp2 hybridization

Three sp2 orbitals on each C atom fuse to form σ bonds with three other C atoms

One 2p orbital remains unhybridized, perpendicular to the plane



Forms of Carbon - Graphite

Structure as per the molecular orbital theory:

All orbitals combine to form MOs

Assist in the stability of graphite

Delocalized electrons account for good electrical conductivity

Used as lubricants in locks

Slipperiness can be attributed to the strong bonding within the C atom layers rather than between the layers

By applying 150,000 atm pressure at 2800°C, graphite can be converted to a diamond



Figure 9.22 (b) - The Structure of Graphite



Figure 9.24 - The p Orbitals and the -Bonding Network in Graphite



Silicon

An important constituent of the compounds that form the earth’s crust

Stable silicon compound involves chains with silicon–oxygen bonds

Fundamental silicon–oxygen compound is silica (SiO2)

Structure

Silicon does not use its 3p orbitals to form strong bonds with oxygen since the orbital size of silicon is larger

Octet rule is satisfied by forming a network of SiO4 tetrahedra with shared oxygen atoms



Structure of Quartz - A Silicate

Quartz holds the same empirical formula as silica, SiO2

It contains chains of SiO4

tetrahedra that share O2

atoms



Silicon and Silicates

Those compounds that are closely related to silica are called silicates

They constitute salts containing metal cations and polyatomic silicon–oxygen anions

Glass is formed when silica is heated above its melting point of 1600°C

Also formed when substances like Na2CO3 are added to the silica melt and then cooled

Properties vary with varying additives



Figure 9.27 - Examples of Silicate Anions



Table 9.5 - Compositions of Some Common Types of Glass



Ceramics

Made from clays, based on silicates, and hardened by firing at high temperatures

Nonmetallic materials that are strong, brittle, and resistant to high temperatures and attack by chemicals

Heterogeneous nature

Formed by the weathering action of water and CO2 on the mineral feldspar, an aluminosilicate

Weathered feldspar produces kaolinite, which consists of small, thin platelets with the empirical formula Al2Si2O5(OH)4



Ceramics

When dry, the platelets stick together, and finally interlock

In the presence of water, they can slide over one another, illustrating the plasticity of clay

After firing:

Remaining water is driven off and silicates and cations form a glass which binds the tiny crystals of kaolinite

Uses:

Drawn into flexible superconducting wires by adding organic polymers

Hold great promise with prosthetic devices



Semiconductors

Electrons of certain elements like silicon can cross the energy gap at 25°C, making them a semiconducting element

Conductivity increases with increasing temperature

n-type semiconductor: A substance whose conductivity is increased by doping with atoms that have more valence electrons than those in the host crystal

p-type semiconductors: A substance whose conductivity is increased by doping with atoms having fewer valence electrons than the atoms of the host crystal



Figure 9.29 - Energy Level Diagrams for: (a) an n-type Semiconductor (b) a p-type Semiconductor



Semiconductors

The connection of a p-type and an n-type form a p–n junctionwhere:

A small number of electrons migrate from the n-type area to the p-type area where low-energy MOs are vacant

Net effect of this migration is to place negative charge on the p-type region and positive charge on the n-type region

This buildup, the contact or junction potential, prevents further migration

Forward bias

Reverse bias



Figure 9.30 (a) - The Charge Carriers of the p-type and n-type Regions

The charge carriers of the p-type region are holes In the n-type region the charge carriers are

electrons



Figure 9.30 (b) and (c) - Reverse and Forward Bias

Section 9.6Molecular Solids


An Introduction to Molecular Solids

Possess strong covalent bonding within molecules but weak forces between molecules

Forces among the molecules depend on the nature of the molecules

CO2, I2, P4, and S8 - No dipole moment; possess London dispersion forces

Molecules that possess dipole moments have greater intermolecular forces, especially when hydrogen bonding is viable

Section 9.7Ionic Solids


Characteristics of Ionic Solids

Stable with high-melting points

Held together by strong electrostatic forces between oppositely charged ions

Structure is based on closest packing of spheres so that:

Electrostatic attraction between oppositely charged ions is maximized

Repulsion among ions with like charges is minimized

For spheres of a given diameter the holes increase in size in the order trigonal < tetrahedral < octahedral



Types of Holes in Closest Packed Structures

Trigonal holes

Formed by three spheres in a given plane

Very small and are never occupied in binary ionic compounds




Tetrahedral holes

Formed when a sphere occupies a dimple formed by three spheres in an adjacent layer

Holes are occupied based on the relative size of the anion and the cation

There are twice as many tetrahedral holes as packed anions in a closest packed structure



Figure 9.33 - The Tetrahedral Hole in the Face-Centered Cubic Unit Cell

(a) The location (red X) of a tetrahedral hole in the face-centered cubic unit cell(b) One of the tetrahedral holes(c) The unit cell for ZnS where the S2– ions (yellow) are closest packed with the Zn2+ ions

(red) in alternating tetrahedral holes




Octahedral holes

Formed by six spheres in two adjacent layers

The number of holes in a ccp structure is the same as the number of packed anions



Interactive Example 9.3 - Determining the Number of Ions in a Unit Cell

Determine the net number of Na+ and Cl– ions in the sodium chloride unit cell

Solution

The Cl–ions are cubic closest packed and thus form a face-centered cubic unit cell

There is a Cl– ion on each corner and one at the center of each face of the cube

Thus the net number of Cl– ions present in a unit cell is:

1 18 + 6 = 4

8 2




The Na+ ions occupy the octahedral holes located in the center of the cube and midway along each edge

The Na+ ion in the center of the cube is contained entirely in the unit cell, whereas those on the edges are shared by four unit cells

Four cubes share a common edge

Since the number of edges in a cube is 12, the net number of Na+ ions present is:

1

1 1 +12 = 44




The net number of ions in a unit cell is 4 Na+ions and 4 Cl– ions, which agrees with the 1:1 stoichiometry of sodium chloride



Table 9.7 - Types and Properties of Solids



Table 9.7 - Types and Properties of Solids (Contd)



Interactive Example 9.4 - Types of Solids

Using the table on the types and properties of solids, classify each of the following substances according to the type of solid it forms

a. Gold

b. Carbon dioxide

c. Lithium fluoride

d. Krypton



Interactive Example 9.4 - Types of Solids

Solution

a. Solid gold is an atomic solid with metallic properties

b. Solid carbon dioxide contains nonpolar carbon dioxide molecules and is a molecular solid

c. Solid lithium fluoride contains Li+ and F– ions and is a binary ionic solid

d. Solid krypton contains krypton atoms that can interact only through London dispersion forces

It is an atomic solid but has properties characteristic of a molecular solid with nonpolar molecules

Section 9.8The Various Forces in Substances: Summary and Comparison


Table 9.8 - lntramolecular (Bonding) Forces



Table 9.8 - lntramolecular (Bonding) Forces (Contd)



Table 9.9 - Intermolecular Forces



Table 9.9 - Intermolecular Forces (Contd)

Section 9.9Vapor Pressure and Change of State


An Introduction

Vaporization, or evaporation is an endothermic process whereby molecules of a liquid escape the liquid’s surface and form gas

Energy required to vaporize 1 mole of a liquid at a pressure of 1 atm is termed the heat of vaporization or the enthalpy of vaporization

Symbolic representation - ΔHvap

Example - Strong hydrogen bonding in the molecules of water in the liquid state leads to a large heat of vaporization



Vapor Pressure

Equilibrium: A point where no further net change can be seen in the amount of liquid or vapor because evaporation and condensation balance each other

The pressure of vapor present at equilibrium is also called equilibrium vapor pressure

Vapor pressure increases with increase in temperature

Sublimation: A substance is said to sublime when it changes directly from solid to gaseous phase without passing through the liquid phase



The Rates of Condensation and Evaporation

The rate of evaporation remains constant and the rate of condensation increases as the number of molecules in the vapor phase increases, until the two rates become equal

At this point, the equilibrium vapor pressure is attained



Figure 9.37 - Estimation of Vapor Pressure Using a Simple Barometer



Concept Check

What is the vapor pressure of water at 100°C?

Justify your answer



Determining Vapor Pressure

Vapor pressure can be determined principally by the size of intermolecular forces in the liquid

Liquids with large intermolecular forces have relatively low vapor pressures

Molecules need a higher energy to escape the vapor phase

Substances with large molar masses have low vapor pressures

Attributed to larger dispersion forces



Vapor Pressure vs. Temperature

Produces a straight line when plotted on a graph

Where

T - Temperature in Kelvin

ΔHvap - Enthalpy of vaporization

R - Universal gas constant

C - Characteristics of a given liquid

ln - Natural logarithm of vapor pressure

vap

vap

Δ 1ln = +

HP C

R T



Figure 9.39(b) - Plots of In(Pvap) Versus 1/T for Water, Ethanol, and Diethyl Ether



The Clausius–Clapeyron Equation

When the values of ΔHvap and Pvap at constant temperature are known, it is easy to calculate the value of Pvap at another temperature

Assume that C does not depend on temperature

1

2

vap, vap

vap, 2 1

Δ 1 1ln =

T

T

P H

P R T T



Interactive Example 9.6 - Calculating Vapor Pressure

The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water at 50°C

Solution

Use the Clausius–Clapeyron equation:

1

2

vap, vap

vap, 2 1

Δ 1 1ln =

T

T

P H

P R T T




For water, we have:

Pvap,T1= 23.8 torr

T1 = 25 + 273 = 298K

T2 = 50 + 273 = 323 K

ΔHvap = 43.9 kJ/mol

R = 8.3145 J/K · mol

Thus,

vap

vap

23.8 torr 43,900 J / mol 1 1ln =

torr 8.3145 J / K mol 323 K 298 K

23.8 ln = 1.37

2

2

,T

,T

P

P




Taking the antilog of both sides,

vap

vap

23.8 = 0.254

= 93.7 torr

2

2

,T

,T

P

P



Changes of State

When a solid is heated, a heating curve can be observed

A plot of temperature against time for a process where energy is added at a constant rate

Enthalpy change that occurs at the melting point is called heat of fusion or enthalpy of fusion (ΔHfus)

Normal melting point: Temperature at which solid and liquid states have identical vapor pressure, and total pressure = 1 atm

Normal boiling point: The temperature a which the vapor pressure of the liquid is exactly 1 atm



Figure 9.41 - Heating Curve for Water



Changes of State

Water can be supercooled and remain in liquid state when:

Temperature < 0°C

Pressure = 1 atm

Liquids can be superheated and raised to temperatures above their boiling point, when heated rapidly The supercooling of water,

represented by S



Concept Check

Which of the following should be larger for a given substance?

Justify your answer

Hvap Hfus

Section 9.10Phase Diagrams


Phase Diagrams - An Introduction

A convenient way of representing the phases of a substance as a function of its temperature and pressure

Triple point: The temperature at which all three phases exist simultaneously

Critical point: The critical pressure and critical temperature, together, define this point

Critical pressure: Pressure required to produce liquefaction at critical temperature

Critical temperature: The temperature above which a liquid cannot be liquefied, irrespective of pressure applied



Figure 9.48 - Phase Diagram for Water

At point X on the phase diagram, water is a solid

As the external pressure is increased while the temperature remains constant, the solid/liquid line is crossed and the ice melts



Phase Diagram for Water - Interpretations and Observations

The solid/liquid boundary line is a negative slope

Melting point of ice decreases with increased external pressure

At the melting point, liquid and solid are in dynamic equilibrium

When pressure is applied, the volume is reduced

A given mass of ice has more volume at 0°C than the same mass of water in liquid state

Freezing point of water is less than 0°C when pressure is greater than 1 atm



Phase Diagram of Water - Applications

Ice skating

Narrow blades of skates exert more pressure

Frictional heat is caused when the skate moves over ice, contributing to further melting of ice

As the blade passes by, the liquid refreezes since normal pressure and temperature are re-established

Low density of ice ensures that the ice formed on rivers and lakes float, ensuring water bodies do not freeze in the winter

In humid climates, snow and ice seem to sublime



Figure 9.49 - Phase Diagram for Carbon Dioxide

The liquid state does not exist at a pressure of 1 atm

The solid/liquid line has a positive slope, since the density of solid carbon dioxide is greater than that of liquid carbon dioxide



Phase Diagram for CO2 - Interpretation and Applications

Interpretations:

Solid/liquid line is a positive slope

Triple point occurs at 5.1 atm and –56.6°C

Critical point can be noticed at 72.8 atm and 31°C

Sublimation occurs at –78°C

Applications:

Dry ice - A convenient refrigerant because it does not undergo the liquid phase under normal atmospheric conditions

Liquid form is used in fire extinguishers at 25°C under high pressure



Concept Check

As intermolecular forces increase, what happens to each of the following?

Boiling point Freezing point

Viscosity Vapor pressure

Surface tension Heat of vaporization

Enthalpy of fusion

chapter 9 liquids and solids · change of state from solid to liquid to gas does not involve change...

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