“Do me a solid..” the story of the other 2 states of matter Copyright©2000 by Houghton Mifflin...

“Do me a solid..”the story of the other 2 states of

matter

Copyright©2000 by Houghton Mifflin Company. All rights

reserved.

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Gas, The oddball of the trio:The Three States of Matter

compared

Solids/Liquids compared to gas

• H2O (s) H2O(l) Δ H = 6 kJ/mol

• H2O (l) H2O(g) Δ H = 41 kJ/mol

• s/l slightly compressible, density similar (see table 10.1)


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Intermolecular Forces

Forces between (not within) molecules in

“condensed states”

dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “” ends attract

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

Figure 10.2Dipole-Dipole Attractions 1% as strong as

the covalent/ionic

bond.

Rapidly weaken with

increase distance


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Figure 10.3 A Water Molecule

H-bonding strong:

small H

big electro. Diff

Water: high melting/

boiling points – to

overcome this force

Also seen with H-F,

H-N bonds


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Figure 10.4The Boiling Points of the Covalent Hydrides of the

Elements in Groups 4A, 5A, 6A, and 7AH-bonding for H-O, H-F, H-N


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What about non-polar molecules and noble gases when condensed? :

London Dispersion Forces


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London Dispersion Forces

relatively weak forces that exist among noble gas atoms and nonpolar molecules when in the condensed state (Ar, C8H18)

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

London Forces

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

• Large molecules, more polarizable.

Explain Table 10.2, pg. 455.

Explain the trend in Halogen b.p.


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Focus on Liquid Properties

Surface Tension

Capillary Action

Viscosity

Surface Tension: The resistance to an increase in its surface area (polar

molecules).

Surface tension explained

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

• Cohesion vs. Adhesion

Viscosity: Resistance to flow (molecules with large intermolecular

forces).

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• High IF, high viscosity (thickness) Ex. Glycerol, gasoline vs. grease

• Ex: add xanthan gum to water, increases the viscosity


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Focus on Solids: 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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Crystalline Solid

Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

Unit Cell: The smallest repeating unit of the lattice. simple cubic

Unit CellsUnit Cells

Unit CellsUnit Cells• Three common types of unit cell.

– Simple cubic, atoms at the corners of a simple cube,• each atom shared by 8 unit cells;

– Body-centered cubic (bcc), atoms at the corners of a cube plus one in the center of the body of the cube,

• corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell;

– Face-centered cubic (fcc), atoms at the corners of a cube plus one atom in the center of each face of the cube,

– corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells.

Unit CellsUnit Cells• Three common types of unit cell.

– Simple cubic, atoms at the corners of a simple cube,• each atom shared by 8 unit cells;

– Body-centered cubic (bcc), atoms at the corners of a cube plus one in the center of the body of the cube,

• corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell;

– Face-centered cubic (fcc), atoms at the corners of a cube plus one atom in the center of each face of the cube,

– corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells.


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Figure 10.9Three Cubic Unit Cells and the Corresponding Lattices

Crystal Structure of Sodium ChlorideCrystal Structure of Sodium Chloride• Face-centered cubic.• Two equivalent ways of defining unit

cell:– Cl- (larger) ions at the corners of the cell, or– Na+ (smaller) ions at the corners of the cell.

• The cation :anion ratio in the unit cell must be a 1:1, since formula is NaCl

• Visualize face center Cl-, Na+ in spaces

Crystal Structure of Sodium ChlorideCrystal Structure of Sodium Chloride

How many Na+, Cl- do you see? Count them!

Body Center = 1 unit

Face Center = 1/2 unit

Edge = 1/4 unit

Corner = 1/8 unit

Three Cubic Unit Cells

and the Correspondin

g Lattices


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How do we determine crystal structure? Interference of x-rays can show us the distance between the particles in the

crystal

the Bragg Equation

• sin θ = xy/d, so xy=d sin θ and yz =d sin θ

• So, xy + yz = d sin θ + d sin θ = 2d sin θ

• For in phase xy + yz = nλ27

The Bragg Equation

• xy + yz = nλ

• xy + yz = 2d sin θ ,

• So, n = 2d sin • d = distance between atoms

n = an integer

= wavelength of the x-rays

See problems 10.1, pg 461 28


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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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Types of Crystalline SolidsIonic 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).

Atomic Solid: atoms at the lattice points. Three kinds of atomic solids: network, metals, group 18. (ex: diamond, Cu, Ar)


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Figure 10.12 Examples of Three Types of Crystalline Solids

Atomic solids compared

• Network, diamond, covalent mp= 3500 C

• Metallic, Cu, delocalized cov., mp = 1083 C

• Group 18, Ar, London, mp = -189 C


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Packing in Metals

• Model: closest packing

• atoms/molecules/ions as spheres.

• To maximize IF forces: sphere packed as

close as possible.

• Small spaces exist between adjacent spheres: called interstitial holes.

Two types of packing possible

hexagonal closest packed (“hcp”)

resulting from aba

cubic closest packed (“ccp”)

resulting from abca


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Let’s pack a crystal hcp and ccpLet’s pack a crystal hcp and ccp 1. Form a base layer of 3 sphere. 2. Place 2nd layer to occupy dimples in layer 1. (one choice only) 3. Two choices for the third layer of

spheres:– 3rd layer over 1st layer(ABAB). Hexagonal

close packing (hcp);– 3rd layer is not over 1st layer. 4th layer is

over 1st layer . (ABCABC. Cubic close packing (ccp).

Choose: which one is ccp, which one is hcp?

Close Packing of SpheresClose Packing of Spheres


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Figure 10.14Hexagonal Closest Packing


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Figure 10.15Cubic Closest Packing

Structures of SolidsStructures of Solids

Close Packing of SpheresClose Packing of Spheres• Each sphere is surrounded by 12 other

spheres (6 in one plane, 3 above and 3 below).

• Coordination number: the number of spheres directly surrounding a central sphere.

• If unequally sized spheres are used, the smaller spheres are placed in the interstitial holes.

The Indicated Sphere Has 12

Nearest Neighbors

True for bothccp and hcp


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Figure 10.15 Cubic Closest Packing


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Figure 10.16The Indicated Sphere Has 12 Nearest Neighbors

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

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Bonding Models for Metals

We can explain why metals are malleable, ductile and conduct heat/electricity

Electron Sea Model: A regular array of metal cations awash 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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Figure 10.18The Electron Sea Model for Metals Postulates

a Regular Array of Cations in a “Sea” of Valence Electrons

Figure 10.19The Molecular Orbital Energy Levels Produced When Various Numbers of Atomic Orbitals Interact.

• 2 ao’s make 2MOs• When 4 ao’s overlap,

4MOs

• Note: when many MO are available, “bands” become virtually

continuous.


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The Band Model for Magnesium

• Core electrons are localized, but valence are in closely spaced Mos: conduction bands.

• Electrons move into empty MOs


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Figure 10.21Two Types of Alloys


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Metal Alloys

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

brass = Cu/Zn

Sterling silver: Ag/Cu

Pewter: tin/Cu/Bi/Sb

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. Ex:

--steel = iron, carbon in the holes

--increase C, reduce ductility, increase hardness and strength.

Mild steel: 0.2% C, nails, chains

Medium: 0.2-0.6% C: beams

High: 1.5% C: springs, tools, knives

Metal alloys

• Alloy steel: mix of interstitial , C and substitutional.

• Ex: stainless steel Fe, C and 11% Cr.

• Cr forms layer of Cr oxide, prevent rust.


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Network Solids

Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.

brittle may or may not conduct heat or electricity carbon, silicon-based

graphite, diamond, ceramics, glass


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Figure 10.22The Structures of Diamond and Graphite: both network solids

Why is diamond insulator not conductor?

a)Carbon in diamond Empty MOs much higher E

b)Metal: Empty MOs close in E

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However, Graphite is a conductor:--each C is sp2, not sp3 like diam.--The p Orbitals in graphite form pi bonds: increase stab. layers & allow conductivity -

Layers in graphite


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Graphite vs. Diamod

• Graphite: conducts, soft, strong bonding in layers, weak between. Layers slide past:

good for lubricant, pencils.

Density: 2.2 g/cc

• Diamond: does not conducts, hardest, strong bonding all around. Good for cutting

Density: 3.5 g/cc

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Mifflin Company. All rights reserved.

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Si: the carbon of geology


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Si, right under C but different

• No long Si-Si chains like C

• Si forms chains of Si-O

• SiO2 = e.f. of silica, most important

• Not CO2 structure, Si cannot

overlap p’s w/ O. No pi bonds.


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• How does Si bond with O?

• A network of SiO4 tetrahedrons.

• No indiv. SiO2, although this is e.f.

• This is quartz


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Glass: melt quartz, cool rapidly: (a) a Quartz Crystal (b) a Quartz Glass

Additives: see table 10.5 Ex: add B2O3 to make borosilcate glass


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Silicate Anions: Note all are based on SiO4 tetrahedron

Ceramics

• Made from clays (contain silicates) that are hardened by firing at high T.

• Non-metallic: silicates bind with cations as clay is dried, trapping crystals of kaolinite.

• Strong, brittle, heat/chemical resistant

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Semiconductors

• Si: Like C, large gap to empty MOs but Si: Like C, large gap to empty MOs but gap is smaller, so some electrons can cross gap is smaller, so some electrons can cross the band gap.the band gap.

• Si is a semiconductor.Si is a semiconductor.

• Conductivity is enhanced by doping with group 13 or group 15 elements.


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Silicon Crystal Doped with (a) Arsenic and (b) Boron

• As: 1 extra valence e-

n-type semiconductor

extra e- can join the conduction band

B: l less valence e-:

p-type semiconductor

creates a “hole” that is filled by an e- from Si.


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Figure 10.30Energy Level Diagrams for (a) an n-Type

Semiconductor and (b) a p-Type Semiconductor

The p-n Junction • e- flow from n to p to

fill “holes”

p side becomes neg• (b) battery with – at

n and + at p. e- flow back: reverse bias, current does not flow

(c) battery other way e- flow right way: forward bias

Application:Transistors:

switches : chips: iphone

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Ionic Solids

• High melting pt• Larger Ions (anions) in

hcp or ccp structure Cations fit in the holes

3 types of holes.

trig< tetra<octa• trig never occupied


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Figure 10.33The Holes that Exist Among Closest Packed Uniform Spheres


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Tetrahedral Holes in a Face-Centered Cubic Unit Cell Ex. ZnS S2- face centered. Zn2+ in tetrahedral holes Q: How many tetrahedral holes? How many are occupied?

NaCl revisited

• Cl- face-centered from ccp packing

• Na+ in all octahedral holes.

• Q: how many octa

holes in (a).


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NaCl octahedral holes

• Note: center Na+

surrounded by the

6 face-centered Cl-

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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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Figure 10.36Behavior of a Liquid in a Closed Container


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Figure 10.37The Rates of Condensation and

Evaporation


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Figure 10.38Vapor Pressure


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Figure 10.39The Number of Molecules in a Liquid With a Given Energy Versus Kinetic

Energy at Two Temperatures


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Figure 10.40The Vapor Pressure of Water


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Figure 10.42 Heating Curve for Water


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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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Figure 10.43The Vapor Pressures of Solid and Liquid Water


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Figure 10.44An Apparatus that Allows Solid and Liquid Water to Interact Only Through the Vapor State


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Figure 10.45Water in a Closed System


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Figure 10.46The Supercooling of Water


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Boiling Point

Constant temperature when added energy is used to vaporize the liquid.

vapor pressure of liquid = pressure of surrounding atmosphere


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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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Figure 10.47The Phase Diagram for Water


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Figure 10.48Diagrams of Various Heating Experiments


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Figure 10.49The Phase Diagram for Water


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Figure 10.52The Phase Diagram for Carbon Dioxide


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Figure 10.50 A Schematic of Two Circuits Connected by a Transistor


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Figure 10.51The Steps for Forming a Transistor in a Crystal of Initially

Pure Silicon

“Do me a solid..” the story of the other 2 states of matter Copyright©2000 by Houghton Mifflin...

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