CONDENSED STATES OF MATTER

Intermolecular forces (IMFs)

States of Matter:

_____ - composed of particles packed closely together with little space between them. Solids maintain a _____ _____ _____.

_____ - any substance that flows. (A fluid) -particles are free to slide past one another and continual change their positions. Particles are in _____ _____ _____ _____ _____

_____ - are fluids composed of particles in _____ _____ _____ _____. Gases are not touching most of the time.

Phase Diagrams

show the phase of matter at a variety of P and T.

Substances can be almost any phase, given the right P and T.

H2O vs. CO2

_____ is less dense in the solid state (ice has lots of space in it), water has a negative slope between the solid & liquid on the phase diagram. Water is densest at 4oC.

_____ goes straight from a solid to a gas at atmospheric pressure.

Intramolecular vs. Intermolecular

_____ __ _____ ___ forces are the forces that hold atoms together in a molecule (within a molecule)

Ex. covalent bonds.

_____ _____ _____ forces are the attractions that molecules have for one another (molecules next to each other)

Ex. H-bonds

Intramolecular vs. Intermolecular

Ex. The forces in H2O

In water, the 2 hydrogens are attached to the oxygen through intramolecular forces.

The reason why water molecules are attached to one another (and not separated like a gas) is because of the intermolecular forces:

The partially negative O of one H2O molecule is attracted to the partially positive H of a different H2O molecule.

Kinetic-Molecular Description of Liquids & Solids

Solids & liquids are _____ _____ _____

atoms, ions, molecules are close to one another

highly incompressible

Liquids & gases are _____

easily flow

_____ _____ _____ _____ in liquids & solids are strong

_____ _____ _____ _____ in gases are weak

IMF types: Ion – ion

_____ _____ attractions are the strongest.

makes up ionic bonding

tend to be crystalline solids (hard, but brittle)

very high melting points

Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-

Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+

Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-Na+Cl-

ion-ion attractions

IMF types: Dipole – Dipole

_____ _____ attractions (fairly strong)

happens with polar molecules b/c they have permanent dipole moments

H—Cl ◦◦◦◦◦◦ H—Cl

δ+ δ- δ+ δ-

IMF types: Hydrogen bonding

_____ _____ _____– a special dipole attraction, stronger than normal dipole-dipole attractions - very strong attraction

~ 2 conditions: must have a _____ _____ bonded to a H and at least one lone pair of electrons on N, O, F

δ+

δ- H . . l

H—N—H ●●●● δ- :N –H δ+

δ+ l δ+ lH Hδ+ δ+

IMF types: Induced Dipole

_____ _____ _____ _____ _____(the weakest), also called _____ _____ _____ _____ attractions

_____ _____ _____ occurs in nonpolar molecules temporary dipole caused by interaction with another

molecule boil and melt very easily Increases with molecular size (aka. The heavier the molecule

the more they exhibit induced dipoles)δ+ δ-

F—F◦ ◦◦ ◦F—F δ- δ+

Mixing of IMF

Ions and molecules can interact with each other in a variety of ways.

ion-dipole

Ions with polar molecules

ion-induced dipole

Ions with nonpolar molecules

dipole-induced dipole

Polar and nonpolar molecules

Mixing of IMF

_____ _____ : An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a molecule that has a dipole.

Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids.

A positive ion (cation) attracts the partially negative end of a neutral polar molecule.

A negative ion (anion) attracts the partially positive end of a neutral polar molecule.

Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases

Ion-Dipole

Ion-Induced Dipole

_____ _____ _____ _____ _____ : An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

Ion-Induced Dipole

Dipole – Induced Dipole

_____ _____ _____ _____ : A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by _____ _____ the arrangement of electrons in the nonpolar species.

Dipole – Induced Dipole

IMF Flow Chart

Properties of liquids

Properties of liquids at constant temperature

~ no definite _____

~ definite _____

~ have surface tension, diffuse, medium density, viscosity, evaporation, capillary action, and vapor pressure. ___ _____ __ of each depends on _____.

_____ ___liquids diffuse into one another

they are ___ ____ in each other for example: water/alcohol gasoline/motor oil

_____ _____ liquids do not diffuse into each other

they are _____ _____ in each other for example: water/oil water/cyclohexane

1. Surface tension

Surface Tension - measure of the _____ _____ _____ _____ that occur at the surface of a liquid

molecules at surface of a liquid are only attracted in a down direction. Denser on top At surface, molecules are attracted downward, thus liquid is

denser on top

water bugs

2. Viscosity

– _____ _____ _____– how easily it flows

~ stronger IMF _____ _____ the liquid is

~ geometry of molecule affects viscosity (more complex shapes = more viscous)

~ very long chains – more viscous b/c longer chains get tangled

C—C—C—C—C—C—C—C —C—C—C—C—C—C—C

more viscous less viscous

3. Capillary action

– tendency of a liquid to be _____ __ or _____ _by a very narrow tube - Stronger IMF more cohesion

~ when a molecule has attraction for itself, it’s called _____ __

~ when a molecule has attraction for other molecules, it is _____ __

capillary rise implies _____ _____ ____ __ (water)

capillary fall implies _____ _____ ____ __ (mercury)

4. Evaporation

– when a liquid changes to the vapor phase at a temp. that is less than it’s boiling point.

Why? If the molecule can gain enough _____, they break through the liquid and go into the atmosphere

5. Vapor pressure

– the pressure of a gas that exists over its solid or liquid state. _____ _____ ____ __

Does not depend on how much liquid/solid you have. Pvap depends on the temp. and type of substance.

You have vapor pressure as long as there is evaporation of a liquid.

Higher Temperature = Higher vapor pressure

Boiling occurs when the Pvap of liquid = Patm

The Liquid State

Vapor Pressure (High _____ = low _____)

Boiling Points & Distillation

~ All liquids have _____ _____ boiling points: based on IMF. The higher IMF , the higher the normal boiling point.

~ can separate liquids on the basis of their b.p. (distillation)

CH3OH has a lower boiling point than C2H5OH, so it changes to a gas first. This means it has lower IMF

Trends in boiling points of Liquids

Gas MW BP(oC)

He 4 -269

Ne 20 -246

Ar 40 -186

Kr 84 -153

Xe 131 -107

Rn 222 -62

The boiling point increases in response to molecular size

The boiling point increases in response to molecular size

Compound MW(amu) B.P.(oC)

CH4 16 -161

C2H6 30 -88

C3H8 44 -42

n-C4H10 58 -0.6

n-C5H12 72 +36

In the Liquid State

Compound MW(amu) B.P.( C)

HF 20 19.5

HCl 37 -85.0

HBr 81 -67.0

HI 128 -34.0

o

HF has the highest B.P. b/c of Hydrogen bonding. The rest increases in response to molecular size.

In the Liquid State

Compound MW(amu) B.P.( C)

H O 18 100

H S 34 - 61

H Se 81 - 42

H Te 130 - 2

o

2

2

2

2

Water has the highest B.P. because of Hydrogen bonding. The rest increases b/c of increase in molecular size.

Various boiling points

Arrange the following substances in order of increasing boiling points.

C2H6, NH3, Ar, NaCl, AsH3

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Amorphous & Crystalline Solids

Amorphous solids _____ ___have a well ordered structure. Particles are irregularly arranged so IMF vary in strength within a sample

Ex. paraffin, glasses

Crystalline solids have well defined structures that consist of extended array of repeating units. Have defined IMF.

give X-ray difraction patterns ~ see Bragg equation in book

Ex. Ice, salt

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Structure of Crystals

unit cell - smallest repeating unit of a crystalEx. bricks are repeating units for buildings

7 basic crystal systems We do not need to learn these 7 – just an FYI for your

future…

Types of Solids

4 Types of _____ _____ _____ _____ :

Covalent

Ionic

Metallic

Molecular

Covalent Solids

Covalent solids: Also known as _____ _____ or _____ _____ _____

individual atoms are covalently bonded to other atoms and those atoms are bonded to other atoms, etc.

In a network solid there are no individual molecules and the entire crystal is considered _____ _____ . This makes covalent solids very hard with very high melting points.

Most are nonconductors.

Formulas for network solids are simple ratios of the component atoms represented by a formula unit (just like ionic compounds)

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Examples of Bonding in Solids

Covalent Solids atoms that are covalently bonded to one another Very strong bonds and hard to break examples:

SiO2 (sand), diamond, graphite, SiC

Ionic Solids

Ionic: positive and negative ions arranged in a specific lattice structure.

_____ _____ _____ _____ are strong.

Also known as a lattice of positive and negative ions held together by electrostatic forces

All of the ionic compounds you are familiar with

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Examples of Bonding in Solids

Ionic Solids ions occupy the unit cell examples:

CsCl, NaCl, ZnS

Metallic Solids

Metallic: metals where each valence electron is thought to belong to the entire structure. It has a closely packed lattice with delocalized electrons throughout.

Metals are seen as a positive nuclei with a _____ of electrons.

The mobility of electrons helps explain the electrical conductivity of metals.

This occurs in all of the metals on the periodic table

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Examples of Bonding in Solids

Metallic Solids positively charged nuclei surrounded by a sea of

electrons

positive ions occupy lattice positions

examples: Na, Li, Au, Ag, ……..

Molecular Solids

Molecular solids: are solids made up of molecules that are next to each other in unit cells held together by the _____ _____ _____ _____.

The attractive forces between individual molecules are _____ _____ _____.

dipole forces are weaker than covalent or ionic bonds, so molecular solids are soft

They are volatile and insulators.

Have relatively low melting temperature

Simple covalent compounds usually form molecular solids

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Examples of Bonding in Solids

Molecular Solids molecules occupy unit cells low melting points,volatile & insulators examples:

water, sugar, carbon dioxide, benzene

Type of Solid

Interactions Properties Examples

CovalentNetwork

Covalent bonding High melting point, hard, nonconducting

C (diamond)SiO2 (quartz)

Ionic Ionic High melting point, conductors in molten or aqueous form, brittle, hard

NaCl, MgO, etc…

Metallic Metallic bonding Conducting, variablehardness and melting point (depends upon strength of metallic bonding)

Fe, Mg, Cu, etc…

Molecular Hydrogen bonding, Dipole-dipole,London Dispersion

Low melting points, (can easily turn into liquids or gases), nonconducting

H2O, CO2

Melt at high temps (most > 1500o C ) because the attractive forces between the individual particles are very strong.

Substance Melting Point (oC)

sand, SiO2 1713

carborundum, SiC ~2700

diamond >3550

graphite 3652-3697

Bonding in Solids - Variations in Melting PointsCovalent Solids

Compound Melting Point (oC)

LiF 842

LiCl 614

LiBr 547

LiI 450

CaF2 1360

CaCl2 772

CaBr2 750

CaI2 740

Bonding in Solids - Variations in Melting PointsIonic Solids

Melt at fairly high temps b/c the attraction between ions are much stronger than in molecular solids but weaker than in covalent solids. Attractive forces

increase as charges on ions increase & their radii decrease. Coulomb’s Law

Melting points vary widely b/c there are large variations in the strengths of metallic bonding. Most metals have fairly high melting points but Mercury is a liquid at room temp.

Metal Melting Point (oC)

Na 98

Pb 328

Al 660

Cu 1083

Fe 1535

W 3410

Bonding in Solids - Variations in Melting PointsMetallic Solids

Have low melting points (most < 300o C ) because the attractive forces between the molecules are rather weak.

Compound Melting Point (oC)

ice 0

ammonia -77.7

benzene, C6H6 5.5

napthalene, C10H8 80.6

benzoic acid, C6H5CO2H 122.4

Bonding in Solids - Variations in Melting PointsMolecular Solids

Alloys are not bonds!

An __________ is a physical mixture made up of two or more metals.

Ex. Brass, Bronze, pewter, steel

Silver and gold used in jewelry is actually an alloy of that element mixed with other metals to make it stronger.

Brief summary

Intermolecular attractions from strongest to weakest

Ion-Ion: ionic compounds (metal/nonmetal)

Hydrogen bonding: H attached to a N, O, or F and lone pair of e- on the central atom

Dipole – Dipole: polar compounds

London Dispersion Forces (induced dipole): all compounds exhibit this, but it is most important with non-polar compounds.

Effects of intermolecular attractions

The compound that has the highest boiling pt, melting pt, and heat of vaporization corresponds to the compound with the strongest IMF.

The highest vapor pressure corresponds to the lowest intermolecular attractions.

If two compounds are nonpolar, the one with the greatest molecular mass has the greater London Forces.

If two compounds are ionic, the one with the greatest charge ions has the greater IMF. If same charge, the smallest ions have the greatest IMF. *Coulomb’s Law

For many years, the world’s record for flying gliders was 60,000 ft. It was set by a Texan who flew into an updraft in front of an approaching storm. The pilot had to fly out of the updraft and land, not because he was out of air (there was still plenty of air in his compressed air bottle) but because he was not wearing a pressurized suit. What would have happened to the pilot’s blood if he had continued to fly higher?