Che108 Chpt 11 Intermolecular Forces

Front 1

Intermolecular forces

Back 1

forces that exist between molecules

weaker than ionic or covalent bonds, therefore, substances can undergo a change of state without a chemical reaction occurring

due to small charges in molecules

Front 2

State Properties

Back 2

Gas
Assumes volume & shape of container, compressible, flows easily, diffuses rapidly

Liquid
Assumes shape of container, does not expand to fill container, virtually incompressible, flows readily, diffuses slowly

Solid
Independent shape & volume, incompressible, does not flow, diffusion occurs slowly

Plasma
high temperature, extreme conditions, e.g. sun

Front 3

Crystalline

Back 3

solids with highly ordered 3d structure

Front 4

State of a substance depends on

Back 4

kinetic energies of particles, interparticle energies of attraction

lesser energies of attraction < gas < liquid < solid < greater energies of attraction

Can be changed by altering temperature or pressure.

Front 5

Boiling Point

Back 5

reflects strength of intermolecular forces

Direct relationship between IMF and boiling point

Boiling: bubbles of a substance's vapor form within the liquid

Front 6

Melting point

Back 6

direct relationship between IMF and melt point of solids

Front 7

Standard Conditions

Back 7

1 atm, 25C

Don't confuse with Standard Temperature & Pressure 1 atm & 0C

Front 8

Normal Boiling & Melting points

Back 8

the temperature of each with P=1 atm

Front 9

Attractive Forces related to Kinetic Energy

Back 9

Attractive > KE: solid
Attractive < KE: gas

Front 10

Strength of IMF

Back 10

Strong: solid
ionic compounds due to strong electrostatic attraction between ions

Moderate: liquid
polar molecular compounds, large molecular compounds due to dipole-dipole attraction, hydrogen bonds, dispersion forces

weak: gas
dispersion forces in small nonpolar molecules

Front 11

Comparison of IMF types

Back 11

Bonding
Ionic: cation:anion, 400-4000 kJ/mol
Covalent: nuclei with shared e-, 150-1100 kJ/mol
Metallic: cation:delocalized e-, 75-1000 kJ/mol

Nonbonding
Ion-dipole: Ion charge to dipole charge, 40-600 kJ/mol
H bond: polar bond to H dipole charge, 10-40 kJ/mol
Dipole-dipole: dipole charges, 5-25 kJ/mol
Ion-induced dipole: ion charge to polarizable e- cloud, 3-15 kJ/mol
Dipole-induced dipole: dipole charge to polarizable e- cloud, 2-10 kJ/mol
Dispersion (London): polarizable e- cloud, 0.05-40 kJ/mol

Front 12

van der Waals forces

Back 12

dipole-dipole attractions, London dispersion forces

Front 13

Dipole

Back 13

a polarized molecule, has a + end and a - end

Front 14

Ion-Dipole Forces

Back 14

between an ion and the partial charge on the end of a dipole

cations to positive end, anions to negative end

Magnitude increases with charge of ion or increasing magnitude of dipole moment

important for dissolution of ionic compounds in polar liquids

Front 15

Dipole-Dipole Forces

Back 15

Neutral polar molecules attract when the + end of one is near the - end of the other

particles must be close, attractive must be stronger than repulsive

particles must be able to rearrange into correct orientation

Liquids: for molecules of ~equal mass & size, strength of IMF increases with increasing polarity

Front 16

Dipole Moment

Back 16

Symbol u
u=Qr
Q is charge in coulombs
r is distance in meters

measure of molecular polarity

Front 17

London Dispersion Forces

Back 17

motion of e- in a nonpolar atom or molecule can create a temporary instantaneous dipole moment, thus causing an attraction

molecules must be very close together

Strength depends on polarizability

Front 18

Polarizability

Back 18

ease with which the e- distribution of an atom or molecule is distorted

Increases with increasing atomic or molecular size, and molecular shape increasingly linear (spherical has less)

Front 19

Hydrogen Bonding

Back 19

intermolecular attraction between H atom in a polar bond (esp. with F, O, or N)) and nonbonding e- pair on a small electronegative ion or atom (usually F, O, or N in another molecule)

Causes the boiling point to skew upward from expected

stabilize protein structure, DNA

responsible for water density being lower as a solid than a liquid

Front 20

Determining IMF

Back 20

All substances have dispersion forces, all attractions are cumulative

weak to strong

No Ions & No polar molecules: dispersion only

No Ions & Yes polar molecules:
Is H bonded to N O or F?
No: dipole-dipole
Yes: hydrogen bond

Yes Ions, Yes polar molecules: ion-dipole

Yes Ions, No polar molecules: ionic bond

Front 21

Viscosity

Back 21

resistance of a liquid to flow

depends on IMF and structural features that allow molecules to tangle

For a series of related compounds, directly related to molecular weight

Front 22

Surface Tension

Back 22

Energy required to increase the surface area of a liquid by a unit amount

water at 20C: 7.29E⁻² J/m²
Hg: 4.6E⁻¹ J/m²
The latter is higher because metallic bonds are stronger than H bonds

Front 23

Cohesive Forces

Back 23

IMF that bind similar molecules

Front 24

Adhesive Forces

Back 24

IMF that bind a substance to a surface

Causes water's concave meniscus

Front 25

Capillary Action

Back 25

the rise of liquids in very narrow tubes

adhesion between liquid & tube walls increases surface area of liquid, surface tension tends to reduce it, liquid is pulled upward

Front 26

Phase Changes

Back 26

changes from one state of matter to another

Gas > Liquid: condensation
Gas > Solid: deposition
Liquid > Solid: freezing
Solid > Liquid: melting
Solid > Gas: sublimation
Liquid > Gas: vaporization

energy of system must change

Front 27

Heat of Fusion

Back 27

ΔHfus

change in enthalpy resulting from the addition or removal of heat from 1 mole of a substance to change its state from a solid to a liquid

Melting point: temperature at which this occurs

Front 28

Heat of Vaporization

Back 28

ΔHvap
energy required to transform a given quantity of a substance into a gas at a given pressure

Boiling point: temperature at which this occurs

Front 29

Vapor Pressure

Back 29

The pressure of a vapor in thermodynamic equilibrium with its condensed phases in a closed system. The gas accumulates above the liquid, exerting pressure on the liquid. All liquids have a tendency to evaporate, and some solids can sublimate into a gaseous form. Vice versa, all gases have a tendency to condense back to their liquid form, or deposit back to solid form.

Inversely related to attractive forces.

Front 30

ΔHfus related to ΔHvap

Back 30

ΔHvap tends to be larger because IMF must be severed

Front 31

ΔHsub

Back 31

heat required to sublime one mole of the substance at a given combination of temperature and pressure

Front 32

Heat of Freezing

Back 32

exothermic to the same degree as ΔHfus

Front 33

Heat of Deposition

Back 33

exothermic to the same degree as ΔHsub

Front 34

Heating Curve

Back 34

Y=temperature, X=amount of heat added

When increasing the temperature to transform a substance from a solid to a liquid and thence to a gas, plateaus occur. Rises are seen within a phase. Plateaus are seen as the substance undergoes state change. The heat is being used to overcome attractive forces between molecules.

Cooling works in an opposite manner.

Front 35

Critical Temperature

Back 35

highest temperature at which a distinct liquid phase can form

Greater than this the substance will be a gas

directly related to IMF

Front 36

Critical Pressure

Back 36

pressure required to bring about liquefaction at critical temperature

directly related to IMF

Front 37

Volatility

Back 37

Directly related to vapor pressure.

Vaporization in an open container, vapor doesn't accumulate to exert pressure. Equilibrium doesn't occur, liquid evaporates entirely.

Front 38

Phase Diagram

Back 38

graphic depiction to summarize conditions of equilibria amongst matter states of a substance

Gas: stable at low pressure, high temp
Solid: stable at high pressure, low temp
Liquid: a curvilinear wedge between the two. Begins at mid-pressure, broadens with increasing temperature and pressure.

Front 39

Triple Point

Back 39

The point in a phase diagram at which all three phases are in equilibrium

Front 40

Crystalline Solid

Back 40

particles ordered in well defined 3d arrays

Typically faces with definite angles, regular shapes.

IMF consistent throughout; therefore, specific melt point.

Front 41

Amorphous Solid

Back 41

particles have no orderly structure

typically large complex molecules

IMF varies throughout sample, melt point varies. These tend to soften over a temp range, then melt.

Front 42

Unit Cell

Back 42

repeating unit of a crystalline solid

7 basic types describe the lattices of all crystalline substances

Front 43

Crystal Lattice

Back 43

3d array of unit cells that represent the crystalline solid

Front 44

Lattice point

Back 44

Each point in a crystal lattice, represents identical environment within the solid

These point denote where atoms are divided between unit cells

Front 45

Cubic Unit Cells

Back 45

Primitive: lattice points at corners only

Body-Centered: lattice points at corners and center of cell (Na)

Face-Centered: lattice points at corners and mid-point of each face (Ni, NaCl)

Front 46

Fraction of an atom that occupies a unit cell

Back 46

Center 1
Face 1/2
Edge 1/4
Corner 1/8

Front 47

Close packing of spheres

Back 47

The particles of crystalline solids are approximately spherical. These arrange to minimize distance between spheres, thus maximizing attractive forces.

Most efficient for equally sized spheres is a hex shape of 7.

Front 48

Hexagonal Close Packing

Back 48

layers of hex shapes stack into the depressions of the hex shape below such that every other layer is identically aligned

Front 49

Cubic Close Packing

Back 49

layers of hex shapes stack into the depressions of the hex shape below such that every third layer is identically aligned

Unit cell is face centered cubic

Front 50

Coordination Number

Back 50

The number of each sphere's equidistant neighbors

Indicates how much of a solid is empty space.

Directly related to packing efficiency.

In hex and cubic close packing, this is 12, spheres are 74%, empty space is 26%

Body-centered: CN=8, spheres 68%

Primitive: CN=6, spheres 52%

With unequal spheres, smaller particles may occupy depressions between larger.

Front 51

Molecular Solids

Back 51

Unit particles: atoms or molecules

Forces: dispersion, dipole-dipole, hydrogen

Properties: soft, poor conductivity, low to moderately high melt point. Dependent upon strength of forces and ability to pack efficiently.

Ex: Ar, methane, sucrose, dry ice

Front 52

Covalent-network Solids

Back 52

Unit particles: atoms connected covalently

Forces: covalent bonds

Properties: very hard, very high melt point, variable conductivity

Ex: diamond (C), quartz (SiO2), graphite (C)

Differences in 3d structure and bonding of carbon yields great differences in physical properties of diamonds vs graphite.

Front 53

Ionic Solids

Back 53

Unit Particles: cation, anion

Forces: electrostatic attraction

Properties: hard, brittle, high melt point, poor conductivity

Ex: salts

Melt point directly related to degree of charge.

Structures dependent on charge & size.

Front 54

Metallic

Back 54

Unit particles: atoms

Forces: metallic bonds

Properties: soft to very hard, low to very high melt point, excellent conductivity, malleable, ductile

Ex: all metallic elements

Structure: typically face centered cubic or body centered cubic with each atom surrounded by 12 or 8 others