8.1 effects of intermolecular forces 8.2 types of intermolecular forces 8.3 liquids 8.4 forces in...

8.1 Effects of Intermolecular Forces

8.2 Types of Intermolecular Forces

8.3 Liquids

8.4 Forces in Solids

8.5 Order in Solids

8.6 Phase Changes

Chapter 8: Effects of Intermolecular Forces

Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.


Learning objective:

Understand the effects of intermolecular forces on condensation, vapourization and melting and boiling points



Intermolecular forces: forces of attraction between molecules which result in liquids and solids.

At STP, only 11 elements are gases.So, intermolecular forces are important in all elements,

except for those which are gases.


Phases of Elements at Room Temperature


The Balance of Kinetic Energy and Intermolecular Potential Energy


Melting and Boiling Points

Both are indicators of the strengths of intermolecular forces: Normal melting point (Tm): the temperature at which a solid

and liquid coexist at equilibrium under a pressure of 1 atm. Normal boiling point (Tb): the temperature at which a liquid

and vapour coexist at equilibrium under a pressure of 1 atm.Vapourization: l → gCondensation: g → l



Learning objectives:

Predict the relative magnitudes of intermolecular forces and their effects on physical properties of substances

Explain the origins of each type of intermolecular force



1. Dispersion Forces – attraction between the negatively charged electron cloud of one molecule and the positively charged nuclei of a neighbor molecule.

2. Dipolar forces – attraction between negatively charged end of one molecule with the positively charged end of another molecule.

3. Induced Dipoles – attraction between an ion or a permanent dipole and a molecule which has had a dipole induced in it by the ion or permanent dipole.

4. Hydrogen bonding forces –attraction between lone pair electrons on an O, N or F atom with a hydrogen atom.


1. Dispersion Forces

Dispersion forces are found in all molecular substancesSuch forces are electrostatic in nature and arise from

attractions involving induced dipoles.Dispersion forces help explain why such things as

nonpolar compounds dissolve in water or ethanol.The magnitude of dispersion forces depends on how

easy it is to polarize the electron cloud of a molecule.A larger molecule has a larger polarizability.


• The process of inducing a dipole is called polarization.

• The higher the molar mass, the higher the polarizability of the molecule.


Polarization

The larger the molecule, the larger the electron cloud.

The larger the electron cloud, the more polarizable the molecule.

The more polarizable, the stronger the dispersion forces.


Larger Molecules are More Polarizable

The larger the atoms or molecules, the stronger the dispersion forces, the higher the boiling point because, larger molecules have higher polarizability and therefore they have stronger dispersion forces.

The stronger forces require more energy to break resulting in a higher boiling point.

Similar arguments for melting points.


Higher Polarization Causes Higher Melting and Boiling Points

Boiling Point vs Molecular Size


Example 8 - 1

Neon and xenon are gases at room temperature, but both become liquids if the temperature is low enough. Draw a molecular picture showing the relative sizes and polarizabilities of atoms of neon and xenon, and use the picture to determine which substance has the lower boiling point.


2. Dipolar Forces

Occur when one polar molecule encounters another polar molecule.

The positive ends will be attracted to the negative ends.

Dipolar forces are characteristically stronger than dispersion forces.

Dipolar forces increase with an increase in the polarity of the molecule.


Acetone is a polar molecule, 2-methylpropane is not.

Therefore, acetone will experience dipolar forces, but 2-methylpropane will not.

Which of the two has a higher boiling point?


2-methylpropane acetone

Example 8 - 2

The line structures of butane, methyl ethyl ether and acetone follow. Explain the trend in boiling points: butane (0°C), methyl ethyl ether (8°C) and acetone (56°C).


3. Hydrogen Bonding Forces

1. One molecule has a hydrogen atom attached by a covalent bond to an atom of oxygen, nitrogen, or fluorine.

2. The other molecule has an oxygen, nitrogen, or fluorine atom.

Remember: they are not really bonds…just attractive forces!


Hydrogen bonding can occur with any hydrogen that is bonded to either oxygen, nitrogen or fluorine.

(N-H, O-H and H-F bonds are very polar)

O H- +

O HH

+

-……

Hydrogen bond

Hydrogen Bonding

H


+

+


Intramolecular H-bond

Hydrogen Bonding Examples

Example 8 - 3

In which of following systems will hydrogen bonding play an important role: CH3F, (CH3)2CO (acetone), CH3OH, and NH3 dissolved in (CH3)2CO?


Binary Hydrogen Compounds


Exceptions due to polarity!

8.3 Liquids

Learning objective:

Explain trends in surface tension, capillary action, viscosity and vapour pressure in terms of intermolecular forces


8.3 Liquids

Liquid Properties

1. Surface Tension2. Capillary Action3. Viscosity4. Vapour Pressure

All are functions of intermolecular forces!


1. Surface Tension

The resistance of a liquid to an increase in its surface area (thus the units J/m2).

The surface molecules of a liquid have a net inward force of attraction, forming a “skin”.

The toughness of the skin is called surface tension.


Capillary Action The upward movement of water inside a capillary

against the force of gravity. Due to large forces between the glass (polar) and the

water (also polar) than among water molecules.


2. Capillary Action

Viscosity - resistance to flow Dependent on intermolecular forces (e.g. hydrocarbons vs.

water) Dependent on length of the carbon chain Dependent on temperature (e.g. viscosity breakdown in

engine oils)


3. Viscosity

Vapour Pressure

Vapour pressure (pvap): the pressure at which dynamic equilibrium is achieved in a closed container.


8.4 Forces in solids

Learning objective:

Explain the properties of solids in terms of the dominant intermolecular forces present


8.4 Forces in solids

In liquids and gases, molecules are free to move continually and randomly.

In solids, molecules, atoms, and ions cannot move freely, but they can vibrate and occasionally rotate.

There are four major solid types: (1) molecular solids, (2) network solids, (3) metallic solids, and (4) ionic solids.


Forces in Solids


1. Molecular Solids

Aggregates of molecules bound together by intermolecular forces.

Molecules of the molecular solids retain their individual properties.

Gases under normal conditions, but form solids at low temperatures.


Crystals of benzoic acid contain pairs of molecules

held together head tohead by hydrogen bonds. These pairs then stack in

planes which are heldtogether by dispersion forces.

The effect of extensive H-bonding is revealed by the high

melting point of glucose (155°C).


2. Network Solids

Unlike molecular solids, network solids have high melting points.

They are held together by covalent bonds which are much harder to break than intermolecular forces.

Bonding patterns determine the properties of the solid.

Usually durable compounds (Rock of Gibraltar, rubies, sapphires, etc…)


Network Solids of Carbon

Diamond – three dimensional array of bonds with each tetrahedral (sp3) carbon linked to others through covalent bonds, this array makes diamonds very strong and abrasive.

Graphite – has trigonal planar (sp2) carbons in a 2D array of bonds, each 2D layer is attracted to its neighbours by dispersion forces.


Network Solids of Carbon


Example 8 - 4

Whereas SiO2 melts at 1710°C, other nonmetal oxides melt at much lower temperatures. For example, P4O6 melts at 25°C. Referring to the accompanying bonding pictures, describe the forces that hold these solids together.


3. Metallic Solids

Bonding of electrons is delocalized.Therefore, the strength of the bonding is variable.Metals are ductile (able to be drawn into wires) and

malleable (able to be hammered into sheets)Metals have a range of melting points


When a metal changes shape, its atoms shift position. However, because the valence electrons are fully delocalized, the energy of these electrons is unaffected.


4. Ionic Solids

Contain cations and anions strongly attracted to each other through interionic forces – in between ions.

Many ionic solids contain metal cations and polyatomic ions

Superconductors: ionic solids composed of oxides of rare earth metals: YBa2Cu3O7-x They carry immense electrical current without

losses due to resistance.


8.5 Order in Solids

Learning objective:

Understand amorphous and crystalline solids at the molecular level


8.5 Order in Solids

Amorphous – No ordered structure to the particles of the solid No well defined faces, angles or shapes Often are mixtures of molecules which do not stack

together well, or large flexible molecules Examples include glass and rubber

Crystalline - The atoms, molecules or ions pack together in an ordered arrangement Such solids typically have flat surfaces, with unique angles

between faces and a unique 3-dimensional shape Examples of crystalline solids include diamonds, and quartz

crystals Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.

The Crystal Lattice and the Unit Cell

Unit cell – the smallest unit from which the entire pattern can be assembled

One unit cell


Basic Definitions

Lattice points – the corners of the unit cell

Crystal lattice – a group of identical unit cells

Cubic Unit Cell – a unit cell which has edges of equal length (l = w = h) and angles of 90°


Three Basic Cubic Crystals

1. Simple cubic (sc) – layers of atoms stacked one directly above another, so that all atoms lie along straight lines at right angles.

2. Body-centered cubic (bcc) – simple cube with one entire atom in the center of the cube (in the body).

3. Face-centered cubic (fcc) – simple cube with atoms in the center of each face of the cube.


Simple Cubic

# 1/8 of an atom in each corner

# atoms = 8 atoms (1/8) = 1 atom


Example 8 - 5

Calculate the mass of a single unit cell of polonium (Po) metal, which crystallizes in a simple cubic structure.


Body Centred Cubic

1/8 of an atom at each corner, and 1 whole atom in the centre

# atoms = 8 atoms (1/8) + 1 = 2 atoms


Face Centred Cubic

1/8 of an atom at each cornerand 1/2 of an atom on each face

# atoms = 8 atoms (1/8) + 6 atoms (1/2) = 4 atoms


Example 8 - 6

Calculate the density of silver metal (in g/cm3), if the edge length of its face-centred cubic unit cell is 407 pm.


Example 8 - 7

Calculate the radius of a palladium (Pd) atom, given that the density of Pd metal is 12.02 g/cm3 and that it has a face-centred cubic unit cell.


Close-Packed Crystals

Close-packing maximizes intermolecular attractions.All empty space around the atoms or molecules is

minimized.


Hexagonal Close-Packed (hcp) and Cubic Close-Packed (ccp)


Ionic Solids

Ions of opposite charges alternate with one another to maximize attractions among ions.

Cations and anions are of different size (cations are usually smaller).


Crystalline Defects

Defects can alter the properties of the solid material

Examine what happens when carbon is added to iron to make steel.

Iron is relatively soft, but adding carbon atoms reduces its ability to become deformed by filling in empty holes in the lattice.


8.6 Phase Changes

Learning objective:

Explain enthalpies of phase changes in terms of intermolecular forces

Interpret a pressure-temperature phase diagram of a pure substance


8.6 Phase Changes

There are three states Solid Liquid Gas

A transformation from one state to another is called a phase change

Each phase change is associated with a change in energy of the system


As a phase change occurs, temperature remains constantChemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.

Temperature and Phase Changes

Heats of Phase Changes

Molar heat of vaporization, Hvap: the heat needed to vapourize one mole of a substance at its normal boiling point.

Molar heat of fusion, Hfus: the heat needed to melt one mole of a substance at its normal melting point.

Molar heat of sublimation, Hsub: the heat needed to sublime one mole of a substance from the solid phase to the gas phase (skips the liquid phase).

Hsub Hvap + Hfus


Example 8 - 8

A swimmer emerging from a pool is covered in a film containing 75 g of water. How much heat must be supplied to evaporate the water?


Phase Diagrams

Illustrate the relationship between phases of matter and the pressure and temperature

The lines identify the conditions under which two phases exist in equilibrium

Triple point – point at which all three phases coexistCritical point – point at which one cannot distinguish

between a gas and a liquid.


A General Phase Diagram


Example 8 - 9

Ammonia is a gas at normal temperature and pressure. Its normal boiling point is 239.8 K, and it freezes at 195.5 K. The triple point for NH3 is p = 0.0610 bar and T = 195.4 K. Use this information to construct an approximate phase diagram for NH3.


Example 8 - 10

A chemist wants to perform a synthesis in a vessel at p = 0.50 atm using liquid NH3 as the solvent. What temperature range would be suitable? When the synthesis is complete, the chemist wants to boil off the solvent without raising the T above 220 K. Is this possible?


Variations on Phase Diagrams

Some substances have more than one solid phase, while others display liquid crystal phases.

The resulting phase diagram must account for all the possible phases.


Silica


Chapter 8 Visual Summary

8.1 effects of intermolecular forces 8.2 types of intermolecular forces 8.3 liquids 8.4 forces in...

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