Chapter 10 Liquids and Solids. Copyright © Cengage Learning. All rights reserved 2 Intramolecular Bonding “Within” the molecule. Molecules are formed

Chapter 10

Liquids and Solids

Copyright © Cengage Learning. All rights reserved 2

Intramolecular Bonding

• “Within” the molecule.• Molecules are formed by sharing electrons between

the atoms.

Intermolecular Forces

Intermolecular forces are attractive forces between molecules.

Intramolecular forces hold atoms together in a molecule.

Intermolecular vs Intramolecular

• 41 kJ to vaporize 1 mole of water (inter)

• 930 kJ to break all O-H bonds in 1 mole of water (intra)

Generally, intermolecular forces are much weaker than intramolecular forces.

“Measure” of intermolecular force

boiling point

melting point

Hvap

Hfus

Hsub

Intermolecular Forces; Explaining Liquid Properties

• The term van der Waals forces is a general term including dipole-dipole and London forces.

Van der Waals forces are the weak attractive forces in a large number of substances.Hydrogen bonding occurs in substances containing hydrogen atoms bonded to certain very electronegative atoms.

Van der Waals Forces 7

• Van der Waals Forces are types of intermolecular forces.

• Consists of:1 Dipole-dipole interactions 2 Ion-dipole interactions3 London forces

Dipole-Dipole Forces

• Polar molecules can attract one another through dipole-dipole forces.

The dipole-dipole force is an attractive intermolecular force resulting from the tendency of polar molecules to align themselves positive end to negative end.

H Cl H Cl


Dipole-Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid


Ion-Dipole Forces

Attractive forces between an ion and a polar molecule

Ion-Dipole Interaction

London Forces

• London forces are the weak attractive forces resulting from instantaneous dipoles that occur due to the distortion of the electron cloud surrounding a molecule.

London forces increase with molecular weight. The larger a molecule, the more easily it can be distorted to give an instantaneous dipole. All covalent molecules exhibit some London force.

• London forces:– exist between all molecules,– is the only attractive force between

nonpolar atoms or molecules,– and dipole-dipole attractions occur

between polar molecules.• Electrons are in constant motion.• Electrons can be, in an instant, arranged in such a way

that they have a dipole. (Instantaneous dipole)• The temporary dipole interacts with other temporary

dipoles to cause attraction.

Intermolecular ForcesDispersion Forces

Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules

ion-induced dipole interaction

dipole-induced dipole interaction

Induced Dipoles Interacting With Each Other

Intermolecular ForcesDispersion Forces Continued

Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.

Polarizability increases with:

• greater number of electrons

• more diffuse electron cloud

Dispersion forces usually increase with molar mass.

SO

O

What type(s) of intermolecular forces exist between each of the following molecules?

HBrHBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.

CH4

CH4 is nonpolar: dispersion forces.

SO2

SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

Hydrogen Bonding• Hydrogen bonding is a force that exists

between a hydrogen atom covalently bonded to a very electronegative atom, X, and a lone pair of electrons on a very electronegative atom, Y.

To exhibit hydrogen bonding, one of the following three structures must be present.

H N OH FH

:::

Only N, O, and F are electronegative enough to leave the hydrogen nucleus exposed.

Hydrogen Bonding

• Hydrogen bonding:

– not considered a Van der Waals Force

– is a special type of dipole-dipole attraction

– is a very strong intermolecular attraction causing higher than expected b.p. and m.p.

• Requirement for hydrogen bonding:

– molecules have hydrogen directly bonded to O, N, or F

Examples of hydrogen bonding:

H2O

NH3

HF

Intermolecular ForcesHydrogen Bond

The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an

electronegative O, N, or F atom.

A H … B A H … Aor

A & B are N, O, or F

Hydrogen Bond

Hydrogen Bonding

• Molecules exhibiting hydrogen bonding have abnormally high boiling points compared to molecules with similar van der Waals forces.

For example, water has the highest boiling point of the Group VI hydrides. Similar trends are seen in the Group V and VII hydrides.

Why is the hydrogen bond considered a “special” dipole-dipole interaction?

Decreasing molar massDecreasing boiling point

Review: Hydrogen Bonding

• A hydrogen atom bonded to an electronegative atom appears to be special.

The electrons in the O-H bond are drawn to the O atom, leaving the dense positive charge of the hydrogen nucleus exposed.It’s the strong attraction of this exposed nucleus for the lone pair on an adjacent molecule that accounts for the strong attraction.A similar mechanism explains the attractions in HF and NH3.

Review: Hydrogen Bonding

H HO

: :

H HO

: :

H HO

: :H H

O: :


Which are stronger, intramolecular bonds or intermolecular forces?

How do you know?

CONCEPT CHECK!CONCEPT CHECK!


Phase Changes

• When a substance changes from solid to liquid to gas, the molecules remain intact.

• The changes in state are due to changes in the forces among molecules rather than in those within the molecules.


Schematic Representations of the Three States of Matter


Phase Changes

• Solid to Liquid– As energy is added, the motions of the molecules

increase, and they eventually achieve the greater movement and disorder characteristic of a liquid.

• Liquid to Gas– As more energy is added, the gaseous state is

eventually reached, with the individual molecules far apart and interacting relatively little.

A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary.

2 Phases

Solid phase - ice

Liquid phase - water


Densities of the Three States of Water


Which molecule is capable of forming stronger intermolecular forces?

N2 H2O

Explain.



Draw two Lewis structures for the formula C2H6O and 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



Which gas would behave more ideally at the same conditions of P and T?

CO or N2

Why?


Liquids

• Low compressibility, lack of rigidity, and high density compared with gases.

• Surface tension – resistance of a liquid to an increase in its surface area:– Liquids with large intermolecular forces tend

to have high surface tensions.


Properties of Liquids

Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area.

Strong intermolecular forces

High surface tension

A molecule within a liquid is pulled in all directions, whereas a molecule on the surface is only pulled to the interior.As a result, there is a tendency for the surface area of the liquid to be minimized.

• Surface tension - a measure of the attractive forces exerted among molecules at the surface of a liquid.

• Surface molecules are surrounded and attracted by fewer liquid molecules than those below.

• Net attractive forces on surface molecules pull them downward.

– Results in “beading”• Surfactant - substance added which decreases the surface

tension

– example: soap

Review: Surface Tension

• Surfactant - substance added which decreases the surface tension

• The nonpolar end is hydrophobic.

• The polar end is hydrophilic.

• Surfactants are amphipathic.

Example - soap

DDBSA is the largest-volume synthetic surfactant because of its relatively low cost, good performance, the fact that it can be dried to a stable powder. DDBSA is also bio-degradable, thus providing environmental friendliness. DDBSA is mainly used to produce household detergents including laundry powders, laundry detergents, dishwashing liquids and other household cleaners. It is used as an emulsifier for applications like agricultural produces and industrial lubricants (Calcium Sulfonate Grease).

DDBSA as a detergent in water:

The Sinking Duck!!!!!!!

Why?

At the air-water interface the hydrophobic end can escape the water, whereas the hydrophilic end remains nestled in the water’s surface. The concentration of water at the interface is reduced. Therefore, the attractive forces between molecules at the interface is reduced, thus the surface tension is reduced.

Liquids

• Capillary action – spontaneous rising of a liquid in a narrow tube:– Cohesive forces – intermolecular forces

among the molecules of the liquid.– Adhesive forces – forces between the liquid

molecules and their container.



Cohesion is the intermolecular attraction between like molecules

Adhesion is an attraction between unlike molecules

Adhesion

Cohesion

• Which force dominates alongside the glass tube – cohesive or adhesive forces?

cohesive forces

Convex Meniscus Formed by Nonpolar Liquid Mercury


Concave Meniscus Formed by Polar Water

• Which force dominates alongside the glass tube – cohesive or adhesive forces?

adhesive forces


Van der Waals Forces and the Properties of Liquids

• Viscosity increases with increasing intermolecular forces because increasing these forces increases the resistance to flow.

Other factors, such as the possibility of molecules tangling together, affect viscosity.Liquids with long molecules that tangle together are expected to have high viscosities.

Intermolecular Forces; Explaining Liquid Properties

• Viscosity is the resistance to flow exhibited by all liquids and gases.

Viscosity can be illustrated by measuring the time required for a steel ball to fall through a column of the liquid. Even without such measurements, you know that syrup has a greater viscosity than water.In comparisons of substances, as intermolecular forces increase, viscosity usually increases.


Viscosity is a measure of a fluid’s resistance to flow.

Strong intermolecular

forces

High viscosity

Kinetic Viscometery

Viscometers

Solids

• Amorphous Solids:– Disorder in the structures– Glass

• Crystalline Solids:– Ordered Structures– Unit Cells


Solid State

• A solid is a nearly incompressible state of matter with a well-defined shape. The units making up the solid are in close contact and in fixed positions.

Solids are characterized by the type of force holding the structural units together.In some cases, these forces are intermolecular, but in others they are chemical bonds (metallic, ionic, or covalent).

A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions.

An amorphous solid does not possess a well-defined arrangement and long-range molecular order.

A unit cell is the basic repeating structural unit of a crystalline solid.

Unit Cell

latticepoint

Unit cells in 3 dimensions

At lattice points:

• Atoms

• Molecules

• Ions

Shared by 8 unit cells

Shared by 2 unit cells

1 atom/unit cell

(8 x 1/8 = 1)

2 atoms/unit cell

(8 x 1/8 + 1 = 2)

4 atoms/unit cell

(8 x 1/8 + 6 x 1/2 = 4)

When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver.

d = mV

V = a3 = (409 pm)3 = 6.83 x 10-23 cm3

4 atoms/unit cell in a face-centered cubic cell

m = 4 Ag atoms107.9 gmole Ag

x1 mole Ag

6.022 x 1023 atomsx = 7.17 x 10-22 g

d = mV

7.17 x 10-22 g6.83 x 10-23 cm3

= = 10.5 g/cm3

Crystal Defects

• There are principally two kinds of defects that occur in crystalline substances.

Chemical impurities, such as in rubies, where the crystal is mainly aluminum oxide with an occasional Al3+ ion replaced with Cr3+, which gives a red color.Defects in the formation of the lattice. Crystal planes may be misaligned, or sites in the crystal lattice may remain vacant.

Bragg Equation

Used to determine the interatomic spacings.

n = integer

= wavelength of the X rays

d = distance between the atoms

= angle of incidence and reflection


= 2 sin n d

Extra distance = BC + CD = 2d sin = n (Bragg Equation)

X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of 14.170. Assuming that n = 1, what is the distance (in pm) between layers in the crystal?

n = 2d sin n = 1 = 14.170 = 0.154 nm = 154 pm

d =n

2sin=

1 x 154 pm

2 x sin14.17= 314.0 pm

Types of Crystalline Solids

• Ionic Solids – ions at the points of the lattice that describes the structure of the solid.

• Molecular Solids – discrete covalently bonded molecules at each of its lattice points.

• Atomic Solids – atoms at the lattice points that describe the structure of the solid.


Examples of Three Types of Crystalline Solids


Classification of Solids


Closest Packing Model

• Closest Packing:– Assumes that metal atoms are uniform, hard

spheres.– Spheres are packed in layers.


The Closest Packing Arrangement of Uniform Spheres• abab 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 so that the spheres in the 3rd layer lie directly over those in the 1st layer.


The Closest Packing Arrangement of Uniform Spheres• abca packing – the spheres in the 3rd layer occupy

dimples in the 2nd layer so that no spheres in the 3rd layer lie above any in the 1st layer.

• The 4th layer is like the 1st.


Hexagonal Closest Packing


Cubic Closest Packing


The Indicated Sphere Has 12 Nearest Neighbors

• Each sphere in both ccp and hcp has 12 equivalent nearest neighbors.


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


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

a) Simple cubicb) Cubic closest packing

a) 1 metal atomb) 4 metal atoms



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.


Length of edge = r 8


Silver metal crystallizes in a cubic closest packed structure. The face centered cubic unit cell edge is 409 pm.

Calculate the density of the silver metal.

Density = 10.5 g/cm3



Bonding Models for Metals

• Electron Sea Model• Band Model (MO Model)


The Electron Sea Model

• A regular array of cations in a “sea” of mobile valence electrons.

Types of Crystals

Metallic Crystals• Lattice points occupied by metal atoms• Held together by metallic bonds• Soft to hard, low to high melting point• Good conductors of heat and electricity• Iron, copper, and silver Cross Section of a Metallic Crystal

nucleus &inner shell e-

mobile “sea”of e-

Band or Molecular Orbital (MO) Model

• Electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal atoms.


Molecular Orbital Energy Levels Produced When

Various Numbers of Atomic Orbitals Interact


The Band Model for Magnesium

• Virtual continuum of levels, called bands.


Metal Alloys

• Substitutional Alloy – some of the host metal atoms are replaced by other metal atoms of similar size.

• Interstitial Alloy – some of the holes in the closest packed metal structure are occupied by small atoms.


Two Types of Alloys

• Brass is a substitutional alloy.

• Steel is an interstitial alloy.


Network Solids


To play movie you must be in Slide Show ModePC Users: Please wait for content to load, then click to play

Mac Users: CLICK HERE

Types of Crystals

diamond graphite

carbonatoms

Covalent Crystals• Lattice points occupied by atoms• Held together by covalent bonds• Hard, high melting point, boiling point• Poor conductor of heat and electricity• Extremely Hard• Diamond

Partial Representation of the Molecular Orbital Energies in a) Diamond b) a Typical Metal


The p Orbitals and Pi-system in Graphite


Ceramics

• Typically made from clays (which contain silicates) and hardened by firing at high temperatures.

• Nonmetallic materials that are strong, brittle, and resistant to heat and attack by chemicals.


Semiconductors

• n-type semiconductor – substance whose conductivity is increased by doping it with atoms having more valence electrons than the atoms in the host crystal.

• p-type semiconductor – substance whose conductivity is increased by doping it with atoms having fewer valence electrons than the atoms of the host crystal.


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

Semiconductor


Silicon Crystal Doped with

(a) Arsenic and (b) Boron







Ionic Solids

• Ionic solids are stable, high melting substances held together by the strong electrostatic forces that exist between oppositely charged ions.


Types of CrystalsIonic Crystals• Lattice points occupied by cations and anions • Held together by electrostatic attraction• Hard, brittle, high melting point, boiling point• Poor conductor of heat and electricity• If dissolves in water, electrolytes• NaCl

CsCl ZnS CaF2

Three Types of Holes in Closest Packed Structures

1) Trigonal holes are formed by three spheres in the same layer.



2) Tetrahedral holes are formed when a sphere sits in the dimple of three spheres in an adjacent layer.



3) Octahedral holes are formed between two sets of three spheres in adjoining layers of the closest packed structures.


• For spheres of a given diameter, the holes increase in size in the order:

trigonal < tetrahedral < octahedral


Types and Properties of Solids


Vapor Pressure

• Liquids are continuously vaporizing.If a liquid is in a closed vessel with space above it, a partial pressure of the vapor state builds up in this space.The vapor pressure of a liquid is the partial pressure of the vapor over the liquid, measured at equilibrium at a given temperature.

Vapor Pressure

• The vapor pressure of a liquid depends on its temperature.

As the temperature increases, the kinetic energy of the molecular motion becomes greater, and vapor pressure increases.Liquids and solids with relatively high vapor pressures at normal temperatures are said to be volatile.

Behavior of a Liquid in a Closed Container a) Initially b) at Equilibrium


Vapor Pressure of a Liquid6

• What happens when you put water in a sealed container?

• Both liquid water and water vapor will exist in the container.

• How does this happen below the boiling point?• Kinetic Theory - Liquid molecules are in continuous

motion, with their average kinetic energy directly proportional to the Kelvin temperature.

The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation

H2O (l) H2O (g)

Rate ofcondensation

Rate ofevaporation=

Dynamic Equilibrium

The Rates of Condensation and Evaporation


Vapor Pressure

• Pressure of the vapor present at equilibrium.• The system is at equilibrium when no net change

occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other.


What is the vapor pressure of water at 100°C? How do you know?

1 atm


Vapor Pressure


Vapor Pressure

• Liquids in which the intermolecular forces are large have relatively low vapor pressures.

• Vapor pressure increases significantly with temperature.


Vapor Pressure vs. Temperature


Clausius–Clapeyron Equation

Pvap = vapor pressure

ΔHvap = enthalpy of vaporization

R = 8.3145 J/K·molT = temperature (in kelvin)


1

2

vap, vap

vap, 2 1

1 1ln =

T

T

P H

P R T T

A Problem to Consider• Carbon disulfide, CS2, has a normal boiling point

of 46oC (vapor pressure = 760 mmHg) and a heat of vaporization of 26.8 kJ/mol. What is the vapor pressure of carbon disulfide at 35oC?

Taking the antiln we obtain:

361)antiln(-0. Hg) mm (760

P 2

Hg mm 760 361)antiln(-0. P2

Hg mm 530 P2

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 65°C.

194 torr


Changes of State




Heating Curve for Water


Which would you predict should be larger for a given substance: ΔHvap or ΔHfus?

Explain why.



• A convenient way of representing the phases of a substance as a function of temperature and pressure:– Triple point– Critical point– Phase equilibrium lines


Phase Diagrams

• A phase diagram is a graphical way to summarize the conditions under which the different states of a substance are stable.

The diagram is divided into three areas representing each state of the substance.The curves separating each area represent the boundaries of phase changes.

Phase Diagrams

• Below is a typical phase diagram. It consists of three curves that divide the diagram into regions labeled “solid, liquid, and gas”. B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• Curve AB, dividing the solid region from the liquid region, represents the conditions under which the solid and liquid are in equilibrium.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• Usually, the melting point is only slightly affected by pressure. For this reason, the melting point curve, AB, is nearly vertical.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• If a liquid is more dense than its solid, the curve leans slightly to the left, causing the melting point to decrease with pressure.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• If a liquid is less dense than its solid, the curve leans slightly to the right, causing the melting point to increase with pressure.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• Curve AC, which divides the liquid region from the gaseous region, represents the boiling points of the liquid for various pressures.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• Curve AD, which divides the solid region from the gaseous region, represents the vapor pressures of the solid at various temperatures.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• The curves intersect at A, the triple point, which is the temperature and pressure where three phases of a substance exist in equilibrium.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• The curves intersect at A, the triple point, which is the temperature and pressure where three phases of a substance exist in equilibrium.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Phase Diagrams

• The temperature above which the liquid state of a substance no longer exists regardless of pressure is called the critical temperature.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.Tcrit

Phase Diagrams

• The vapor pressure at the critical temperature is called the critical pressure. Note that curve AC ends at the critical point, C.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.Tcrit

Pcrit

Phase Diagram for Carbon

Dioxide


Phase Diagram for

Water


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

– Boiling point– Viscosity– Surface tension– Enthalpy of fusion– Freezing point– Vapor pressure– Heat of vaporization


Documents

Chapter 10 Liquids and Solids. Copyright © Cengage Learning. All rights reserved 2 Intramolecular Bonding “Within” the molecule. Molecules are formed