Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Liquids and Solids Chapter 10.

Copyright©2000 by Houghton Mifflin Company. All rights reserved.

1

Liquids and Solids

Chapter 10


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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


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• H2O(s) → H2O (liq) ΔHfus = 6.02 kJ/mol

• H2O(liq) → H2O (g) ΔHvap = 40.7 kJ/mol


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Schematic representations of the three states of matter.


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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


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Forces between (rather than within) molecules.

dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “” ends of the dipoles are close to each other.

hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)


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Dipole-Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid


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(a) The electrostatic interaction of two polar molecules. (b) The interaction of many dipoles in a condensed state.


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(a) The polar water molecule. (b) Hydrogen bonding among water

molecules.


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

Decreasing molar massDecreasing boiling point

Why this sudden increase?

The boiling point represents the magnitude and type of bonding


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Ion-Dipole Forces

Attractive forces between an ion and a polar molecule

Ion-Dipole Interaction


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London Dispersion Forces

relatively weak forces that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18)

caused by instantaneous dipole, in which electron distribution becomes asymmetrical.

the ease with which electron “cloud” of an atom can be distorted is called polarizability.


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(a) An instantaneous polarization can occur on atom A, creating an instantaneous dipole. This dipole creates an induced dipole on neighboring atom B. (b) Nonpolar molecules such as H2 also can develop instantaneous and induced dipoles.


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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.


17

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.


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Properties of Liquids

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

Or

The resistance to an increase in its surface area (polar molecules).

Strong intermolecular

forces

High surface tension


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A molecule in the interior of a liquid is attracted by the molecules surrounding it, whereas a molecule at the

surface of a liquid is attracted only by molecules below it and on each side.


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More Properties of Liquids

Capillary Action: Spontaneous rising of a liquid in a narrow tube.

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


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Cohesion is the intermolecular attraction between like molecules

Adhesion is an attraction between unlike molecules

Adhesion

Cohesion


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

Strong intermolecular

forces

High viscosity


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Types of Solids

Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)].

Amorphous solids: considerable disorder in their structures (glass).


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

A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing

Crystallinequartz (SiO2)

Non-crystallinequartz glass


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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 (smallest) repeating structural unit of a crystalline solid.

Unit Cell

latticepoint

At lattice points:

• Atoms

• Molecules

• Ions

Unit cells in 3 dimensions


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Shared by 8 unit cells

Shared by 2 unit cells


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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)


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Three cubic unit cells and the corresponding lattices.


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Analysis of crystal structure using X-Ray Diffraction


34Extra distance = BC + CD = 2d sin = n (Bragg Equation)


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Bragg Equation

Used for analysis of crystal structures.

n = 2d sin

d = distance between atoms

n = an integer

= wavelength of the x-rays


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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= 77.0 pm

11.5


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Types of Crystalline Solids

Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).


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Types of Crystals

Ionic Crystals• Lattice points occupied by cations and anions• Held together by electrostatic attraction• Hard, brittle, high melting point• Poor conductor of heat and electricity

CsCl ZnS CaF2


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S.Ex 10.2• Silver crystallizes in FCC closest packed structure.

The radius of silver atom is 145 pm. Calculate the density of solid silver.

• d2 + d2 = (4r)2

2d2 =16 r2

=16x1452

d = 409 pm


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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


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Types of Crystals

Molecular Crystals• Lattice points occupied by molecules• Held together by intermolecular forces• Soft, low melting point• Poor conductor of heat and electricity


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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

Cross Section of a Metallic Crystal

nucleus &inner shell e-

mobile “sea”of e-


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Examples of three types of crystalline solids. (a) An atomic solid. (b) An ionic

solid. (c) A molecular solid.


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Types of Crystals


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Bonding Models for Metals

Electron Sea Model: A regular array of metals in a “sea” of electrons.

Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.


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The electron sea model for metals postulates a regular array of cations in a "sea" of valence electrons. (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.


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The molecular orbital energy levels produced when various numbers of atomic orbitals interact.


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(left) A representation of the energy levels (bands) in a magnesium crystal. (right) Crystals

of magnesium grown from a vapor.


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Metal Alloys

1. Substitutional Alloy: some metal atoms replaced by others of similar size.

brass = Cu/Zn

Substances that have a mixture of elements and Substances that have a mixture of elements and metallic properties.metallic properties.


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Metal Alloys(continued)

2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are

occupied by small atoms.

steel = iron + carbon3. Both types: Alloy steels contain a

mix of substitutional (carbon) and interstitial (Cr, Mo) alloys.


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Two types of alloys: (a) substitutional (b) interstitial


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A “steaming” piece of dry iceCO2


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(a) Sulfur crystals (yellow) contain S8 molecules. (b) White phosphorus (containing P4

molecules) is so reactive with the oxygen in air that it must be stored under water.


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Ionic Solids

• Ionic solids are stable, high melting point and held ions together by stron electrostatic forces between opposite ions.

• Structure of most binary ionic solids (NaCl) can be explained by the closest packing of spheres– Anion (large ion) packed in one of the closest packing

(hcp or ccp)

– Cation (small ion) fits in holes among the anions.


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(a) The locations (gray X) of the octahedral holes in the face-centered cubic unit cell. (b) Representation of the unit cell for solid NaCl.


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S.Ex 10.3

• Determine the net number of Na+ and Cl- ions in the sodium Chloride unit cell.

• Cl- form closest packed FCC– 8(1/8) + 6(1/2) = 4

Na+ one in the center of cube, (not shared)

Ones in the edge are shared by 4 unit cells; there are 12 edges;

1 + 12(1/4) = 4 1:1 stoichiometry


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Vapor Pressure

. . . is the pressure of the vapor present at equilibrium.

. . . is determined principally by the size of the intermolecular forces in the liquid.

. . . increases significantly with temperature.

Volatile liquids have high vapor pressures.


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Behavior of a liquid in a closed container.


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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


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BeforeEvaporation

At Equilibrium


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(a) The vapor pressure of a liquid can be measured easily using a simple barometer of the type shown here. (b) The three liquids, water, ethanol (C2H5OH), and diethyl ether [(C2H5)2O], have quite different vapor pressures.


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(a) The vapor pressure of water, ethanol, and diethyl ether as a function of temperature. (b) Plots of In(Pvap) versus 1/T (Kelvin temperature) for water, ethanol, and diethyl ether.


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Molar heat of vaporization (Hvap) is the energy required to vaporize 1 mole of a liquid.

ln P = -Hvap

RT+ C

Clausius-Clapeyron EquationP = (equilibrium) vapor pressure

T = temperature (K)

R = gas constant (8.314 J/K•mol)


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LnPT1

PT2

=Hvap

R

1 1

T2 T1


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S.Ex 10.7 • The vapor pressure of water at 25oC is 23.8 torr, and the

heat of vaporization of water at 25 oC is 43.9 kJ/mol. Calculate vapor pressure of water at 50 oC.


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The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure.

The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.


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Melting Point

Molecules break loose from lattice points and solid changes to liquid. (Temperature is constant as melting occurs.)

vapor pressure of solid = vapor pressure of liquid


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Mel

ting

11.8F

reez

ing

H2O (s) H2O (l)

The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium


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Molar heat of fusion (Hfus) is the energy required to melt 1 mole of a solid substance.


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The supercooling of water. The extent of supercooling is given by

S.


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Sub

limat

ion

Dep

ositi

on

H2O (s) H2O (g)

Hsub = Hfus + Hvap

( Hess’s Law)

Molar heat of sublimation (Hsub) is the energy required to sublime 1 mole of a solid.


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Phase Diagram

Represents phases as a function of temperature and pressure.

critical temperature: temperature above which the vapor can not be liquefied.

critical pressure: pressure required to liquefy AT the critical temperature.

critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).


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The phase diagram for water.


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The critical temperature (Tc) is the temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure.

The critical pressure (Pc) is the minimum pressure that must be applied to bring about liquefaction at the critical temperature.


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Diagrams of various heating experiments on samples of water in a closed system.

Negative Slope


78

The phase diagram for carbon dioxide.

Positive Slope

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Liquids and Solids Chapter 10.

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