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Chapter 11 : Liquids, Solids, and Intermolecular Forces. Outline Climbing Geckos !?! Solids, Liquids, and Gases Intermolecular forces Surface tension, viscosity and capillary action Water Crystalline Solids. Climbing Geckos.

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slide1

Chapter 11 : Liquids, Solids, and Intermolecular Forces

  • Outline
  • Climbing Geckos !?!
  • Solids, Liquids, and Gases
  • Intermolecular forces
  • Surface tension, viscosity and capillary action
  • Water
  • Crystalline Solids
slide2

Climbing Geckos

A gecko can adhere to a glass ceiling with a single toe ! Given that an average, full grown gecko weighs 100g calculate the force of adhesion it must apply to the glass to do so ?

Climbing Geckos

slide3

Solids, Liquids, and Gases

Solids, Liquids, and Gases

slide4

Solids

  • Compared to gases, solids are dense
  • Have a definite shape and volume

Liquids

  • Compared to gases, solids are dense
  • Have an indefinite shape and volume

Solids, Liquids, and Gases

slide5

Phase Changes

Liquid propane is a good example of a pressure-induced phase change.

Solids, Liquids, and Gases

slide6

Phase Changes

Remember : A phase change is a physical change.

If this picture represents H2O in the liquid state….

Solids, Liquids, and Gases

slide7

When a molecule in a liquid state undergoes a phase change to a gas, it must break all the intermolecular forces acting upon it.

Solids, Liquids, and Gases

slide8

Intermolecular Forces

Intramolecular forces are interactions of atoms within a molecule.

Intermolecular forces are interactions between molecules.

Among other things, intermolecular forces are responsible for condensation of gases, and freezing of liquids into solids.

Examples of intermolecular forces in liquids;

  • Surface tension
  • Viscosity
  • Vapour pressure

Intermolecular Forces

slide9

δ-

δ-

δ+

δ+

Coulomb’s Law

We’ve seen this before when we discussed bonding.

In the case of bonding, q = 1.6 x 10-23 C.

But we know that molecules can have partial charges if they have polar bonds.

The interactions between these partial charges are the basis of intermolecular forces.

Intermolecular Forces

slide10

Because we are not talking about full charges, q < 1.6 x 10-23 C.

What does this tell you about the strength of intermolecular forces compared to bonding forces ?

Intermolecular forces also act over much larger distances compared to electron-proton interactions.

Intermolecular Forces

slide11

Types of IM Forces

There are four categories of IM forces that we will consider.

  • Dispersion Forces
  • Dipole-Dipole Forces
  • Hydrogen Bonding
  • Ion-Dipole Forces

Intermolecular Forces

slide12

Dispersion (aka London) Forces

He is a noble gas and exists as a free atom in the gas state at STP.

If there are no bonds, there can be no permanent dipoles. Therefore no partial charges.

If partial charges lead to IM forces and IM forces are needed to condense gases into liquids, does that mean you can’t liquefy Helium ?

Intermolecular Forces

slide13

At any instant, the electrons of He (or any other atom) may not be symmetrically distributed about the nucleus.

At such an instant, the left side of He will have a slight negative charge whereas the right side will have a slight positive charge.

Intermolecular Forces

slide14

Instantaneous dipoles create induced dipoles in nearby atoms.

The forces that pull molecules together due to instantaneous and induced dipoles are called dispersion or London forces.

The magnitude of dispersion forces depends on the ability of an atom’s (or molecule’s) electron cloud to be distorted.

nucleus

Intermolecular Forces

slide15

Simple Example

Match the following noble gases and their boiling points.

He

4.2 K

Ar

165 K

Xe

87 K

All molecules are comprised of electrons therefore ALL molecules have dispersion forces.

Intermolecular Forces

slide16

So is it as simple as bigger atoms (or molecules) have larger dispersion forces ?

The elongated shape of n-pentane means there is more area of the molecule available to interact with other molecules.

Intermolecular Forces

slide18

Dipole-Dipole Forces

Recall our discussion on polar covalent bonds.

These partial charges (resulting from the permanent dipole moment) cause the molecules to order themselves.

This is an electrostatic interaction. The larger the dipole moment, the larger the partial charges and the greater the dipole-dipole interaction.

Intermolecular Forces

slide19

Which of the following molecules would be expected to have the larger dipole-dipole forces ?

Intermolecular Forces

slide21

Example

Which of the following molecules will have dipole-dipole forces ?

  • CO2
  • CH2Cl2
  • CH4

Intermolecular Forces

slide22

Hydrogen Bonding

We expect b.p. to increase with increasing molecular mass. What’s going on with HF, NH3, and H2O ?

Intermolecular Forces

slide23

H

F

When hydrogen is covalently bonded to a very electronegative atom an extremely polar bond forms.

Consequently, the hydrogen atom is simultaneously attracted to the electronegative atom (O,N,F) of a neighbouring molecule.

An example of a hydrogen bonding network.

In the gas state, HF exists in a cyclic “supramolecular” structure consisting of six HF molecules linked with h-bonds.

Intermolecular Forces

slide24

Hydrogen Bonding in Water

Crystal structure of ice

Four h-bonds per molecules

“Open” structure

Structure of liquid water

Some h-bonds broken

“Denser” structure

Intermolecular Forces

slide25

Why Ice Floats

Intermolecular Forces

slide26

Intramolecular and Intermolecular H-Bonding

Intermolecular h-bonding occurs between molecules.

Intramolecular h-bonding occurs between groups on the same molecules.

Intermolecular Forces

slide27

How Strong are Intermolecular Forces ?

Ionic bond

Covalent bond

Ion-Dipole

Hydrogen bonding

Perm. Dipole – Perm. Dipole

Induced dipole forces

Intermolecular Forces

slide28

Examples

2. Arrange the following in order of increasing boiling point:

He, Ne, O3, O2, Cl2, (CH3)2C=O

3. Arrange the following in order of increasing boiling point:

C6H6 ; C6H5Cl ; C6H5Br ; C6H5OH

Intermolecular Forces

slide29

Surface Tension

Molecules at the surface of a liquid

Molecules in the interiour of a liquid

Because of the increased number of interactions, molecules in the bulk of solution are at a lower energy state than those on the surface.

Surface Tension, Viscosity, Capillary Action

slide30

To increase the area of the interface, molecules most be moved from the interiour of the liquid to the interface. This requires work be done…NOT energetically favourable.

Surface tension is the work (or energy) required to increase the area of a liquid

Mathematically :

Why are raindrops spherical and not cubic in shape ?

Surface Tension, Viscosity, Capillary Action

slide31

Another example. How does a steel nail float on water ?

Surface Tension, Viscosity, Capillary Action

slide32

Wetting

Cohesive forces are intermolecular forces between molecules.

Adhesive forces are intermolecular forces between molecules.

When a liquid spreads into a film on a surface we say the liquid wets the surface. This will happen if the A.F. > C.F.

Surface Tension, Viscosity, Capillary Action

slide33

Meniscus formation is another example of wetting phenomena.

Surface Tension, Viscosity, Capillary Action

slide34

Viscosity

Viscosity describes resistance to flow. Liquids such as honey and ethylene glycol have high viscosity while ethanol and water have low resistance to flow.

Strong intermolecular forces cause high viscosities.

Unit of viscosity is the Poise, (P), = g cm-1 s-1

Surface Tension, Viscosity, Capillary Action

slide35

As temperature increases, the intermolecular forces are less pronounced and the viscosity decreases.

But shape matters too !

Surface Tension, Viscosity, Capillary Action

slide36

Vaporization and Vapor Pressure

  • molecules in the liquid are constantly in motion
  • Avg. KE is proportional to temperature
  • But the KE is distributed over a range of values

Surface Tension, Viscosity, Capillary Action

slide37

if these molecules are at the surface, they can escape from the liquid and become a vapor

  • At a given temperature, the larger the surface area, the faster the rate of evaporation
  • As the temperature increases, a greater fraction of water molecules have this minimum energy  faster rate of evaporation
  • The stronger the IMFs, the slower the rate of vaporization
slide38

What happens to the molecules that have escaped ?

  • some molecules of the vapor will lose energy through molecular collisions
  • the result will be that some of the molecules will get captured back into the liquid when they collide with it
  • also some may stick and gather together to form droplets of liquid
    • particularly on surrounding surfaces
slide39

Evaporation vs. Condensation

  • vaporization and condensation are opposite processes
  • in an open container, the vapor molecules generally spread out faster than they can condense. Net result is a loss of liquid.
  • however, in a closed container, the vapor is not allowed to spread out indefinitely
  • the net result in a closed container is that at some time the rates of vaporization and condensation will be equal
slide40

Energetics of Vaporization

  • when the high energy molecules are lost from the liquid, it lowers the average kinetic energy
  • if energy is not drawn back into the liquid, its temperature will decrease – therefore,
  • vaporization requires input of energy to overcome the attractions between molecules
slide41

Heat of Vaporization

  • the amount of heat energy required to vaporize one mole of the liquid is called the Heat of Vaporization, DHvap
    • sometimes called the enthalpy of vaporization
  • always endothermic, therefore DHvap is +
  • somewhat temperature dependent
slide42

Dynamic Equilibrium

Dynamic Equilibrium is when the rate of vaporization equals the rate of condensation

slide43

the pressure exerted by the vapor when it is in dynamic equilibrium with its liquid is called the vapor pressure

  • the weaker the attractive forces between the molecules, the more molecules will be in the vapor
    • the higher the vapor pressure, the more volatile the liquid
slide44

Boiling Point

  • when the temperature of a liquid reaches a point where its vapor pressure is the same as the external pressure, vapor bubbles can form anywhere in the liquid
  • this phenomenon is what is called boiling and the temperature required to have the vapor pressure = external pressure is the boiling point

At the summit of Mt. Everest, water boils at 78oC. Why ?

slide45

Heating Curve of a Liquid

  • as you heat a liquid, its temperature increases linearly until it reaches the boiling point
    • q = mass xCsxDT
  • once the temperature reaches the boiling point, all the added heat goes into boiling the liquid – the temperature stays constant
slide48

Melting = Fusion

  • as a solid is heated, its temperature rises and the molecules vibrate more vigorously
  • once the temperature reaches the melting point, the molecules have sufficient energy to overcome some of the attractions that hold them in position and the solid melts (or fuses)
  • the opposite of melting is freezing
slide49

Energetics of Melting

  • when the high energy molecules are lost from the solid, it lowers the average kinetic energy
  • if energy is not drawn back into the solid its temperature will decrease
    • and freezing is an exothermic process
  • melting requires input of energy to overcome the attractions between molecules
  • the amount of heat energy required to melt one mole of the solid is called the
    • sometimes called the enthalpy of fusion
slide50

Sublimation and Deposition

  • Similar to the process in a liquid, molecules in a solid can break free from the surface and become a gas – this process is called sublimation
  • the capturing of vapor molecules into a solid is called deposition
  • the solid and vapor phases exist in dynamic equilibrium in a closed container
    • at temperatures below the melting point
    • therefore,molecular solids have a vapor pressure
slide51

Supercritical Fluid

  • as a liquid is heated in a sealed container, more vapor collects causing the pressure inside the container to rise
    • the density of the vapor increases
    • the density of the liquid to decreases
  • at some temperature, the meniscus between the liquid and vapor disappears and the states commingle to form a supercritical fluid
  • supercritical fluid have properties of both gas and liquid states
slide52

The Critical Point

  • the temperature required to produce a supercritical fluid is called the critical temperature
  • the pressure at the critical temperature is called the critical pressure
  • at the critical temperature or higher temperatures, the gas cannot be condensed to a liquid, no matter how high the pressure gets