chemistry 100 – chapter 11 intermolecular forces

Chemistry 100 – Chapter 11

Intermolecular Forces

Intramolecular and Intermolecular Forces We have just been discussing the covalent

bond - the force that holds atoms together making molecules.

We have also talked about the ionic bond. These are intramolecular forces.

There are also forces that cause molecules to attract each other. These are called intermolecular forces.

Intermolecular Forces

Salt (NaCl) is a solid because of the strong electrostatic attraction of the Na+ and the Cl–

ions: the ionic bond Q: Why is Cl2 a gas, Br2 a liquid and I2 a

solid? A: Intermolecular forces Q: Why is molasses “thick” while water has

low viscosity? Same answer. Q: Why is it possible to float a needle on

water?

Intermolecular Forces (cont’d)

Q: Why is water a liquid but H2S is a gas at 25ºC?

Q: Why are real gases not ideal? A: van der Waals forces (same thing, different name)

States of Matter

Gas: No defined shape or volume, compressible, rapid diffusion, flows readily

Liquid: Takes shape of container, virtually incompressible, low diffusion, flows readily

Solid: Has its own shape and volume, virtually incompressible, extremely slow diffusion, does not flow

Liquids and solids are called condensed phases because particles are close together

Intermolecular Forces

Much weaker than chemical bonds. Covalent bonds 200 kJ/mol and more.

Intermolecular forces less than 50 kJ/mol

When a liquid vaporizes, the intermolecular forces must be overcome. But no covalent bonds are broken.

Ion - dipole forces

Dissolve an ionic compound in water

Charged ions interact with the dipole of water molecules

Dipole-dipole Forces

Interactions between the dipoles in a polar liquid leads to a net attraction

Dipole - Induced dipole force

In a mixture of two liquids were one is polar and the other is not, the dipole of the polar molecule can induce a dipole in the other.

The energy of this force depends on the polarizabilty of the non-polar molecule. Larger molecules are more polarizable

London Dispersion Forces

How do we account for the fact that non-polar gases can be liquefied and solidified?

Fritz London proposed instant dipoles. These are dipoles that result from the random movement of the electron cloud.

Dispersion forces work over very short distances

Polarizability

The ease with which a dipole can be induced depends on the polarizability of the molecule

Large molecules are more polarizable - easier to distort the electron cloud

Polarizability and Boiling Points

Boiling points (K) of halogensF2 Cl2 Br2 I285.1 238.6 332.0 457.6

Boiling points of Noble gasesHe Ne Ar Kr Xe4.6 27.3 87.5 120.9

166.1

Dipole or Dispersion? Dispersion forces operate between all

molecules - polar and non-polar Molecules with comparable molecular

weights and shapes have approximately equal dispersion forces. Any difference is due to dipole-dipole

attractions When molecules differ widely in

molecular weight, dispersion forces tend to be the decisive ones

Water And Ammonia Have Unusual Boiling Points!!

The Hydrogen Bomb (err – Bond)

A special dipole-dipole intermolecular attraction H in a polar bond (H-F, H-O or H-N) an unshared electron pair on a nearby electronegative

ion or atom (generally F, O or N) Hydrogen bonds (4 to 25 kJ/mol, or larger) are

weaker than covalent bonds but stronger than most dipole-dipole or dispersion forces.

Molecules with H bonding

HF (can behave as if it were H2F2) NH3

H2O (ice is less dense than liquid at 0ºC) alcohols (e.g. CH3OH) amines (e.g CH3NH2) carboxylic acids (e.g. CH3COOH)

Summary

Liquids - viscosity

Viscosity (resistance to flow) Molecules that have strong intermolecular

forces cannot move very easily - more viscous

Viscosity decreases at higher temperatures. The kinetic energy overcomes the intermolecular forces

Liquids - surface tension Surface tension: the energy that

must be expended to increase the surface of a liquid

Surface tension

Liquids with strong intermolecular bonds have high surface tensionsWater: strong H-bonds, so high surface tension, 7.29 10–2 J/m2

Mercury: atoms held by metallic bonds, so even higher surface tension, 4.6 10–1 J/m2

Wetting & capillary action

Water wets clean glass (spreads out) but beads on a waxy surface

Water climbs up a capillary tube Mercury does not wet glass Mercury level is depressed in

capillary tube

To wet or not to wet?

Competition between two tendencies Cohesive forces; intermolecular forces

that bind similar molecules together. Keeps liquid as a bead

Adhesive forces: Intermolecular forces between molecules of a liquid and those of a surface. Makes liquid spread out

Are we all wet?? Water on clean glass

adhesive forces > cohesive forcesexplains the upward meniscus of water

Water on polished tablecohesive forces win because H2O molecules are not attracted to wax

Mercury on glass - cohesive forces win because Hg molecules are not attracted to glass

Explains depression in capillary tube and downward meniscus

Phase changes

Phase changes

Enthalpy of fusion or heat of fusion for water Hfus = 6.01 kJ/mol.

Enthalpy of vaporization or heat of vaporization for water Hvap = 40.67 kJ/mol. cooling effect of evaporation refrigeration (Not Freon-12) steam burn generally severe

Heating Curve

Supercooling, superheating

A liquid cooled below its freezing point is said to be supercooled requires very clean conditions the molecules are moving slowly but have not

organized themselves into the solid form A liquid heated above its boiling point is

said to be superheated a danger with heating water in microwave oven

Critical T and P

A gas can be liquefied by cooling A gas can be liquefied by increasing pressure but

only if the temperature is below the compound’s critical temperaturesubstance critical temp K and Cammonia 405.6 (133)carbon dioxide 304 (31)argon 150.9 (-122)

Critical pressure: pressure needed to liquefy gas at critical temperature

Substance at Tc and Pc - supercritical fluid

Vapour Pressure

Every liquid in a closed container gives off vapour until a certain pressure is reached - the liquid’s vapour pressure.

The vapour pressure of a liquid increases with increase in temperature.

We can explain these facts using the kinetic theory

Volatile

A liquid in an open container will evaporate

As vapour moves away, the liquid releases more molecules into the vapour phase to try to build up to the correct vapour pressure

Liquids with high vapour pressure evaporate more quickly - they are volatile

Boiling point: the temperature at which vp = 760 torr

Boiling

A liquid boils when the vp equals the atmospheric pressure

Normal boiling point temp -> when vp is 760 torr. We list normal bp values in textbooks

Actual boiling point -> then the liquid has a vp equal to the external atmospheric pressure

Water boils at temperature lower than 100ºC atop mountains - it never reaches 100ºC

Water boils at higher temp in pressure cooker

Phase diagrams

A graphical way to show the equilibria between different phases of a substance

Thing to look for: critical point triple point ( three phases) how bp (and mp) varies with pressure

Triple point is not pressure dependent - the vapour has to be at the critical pressure! useful for thermometer calibration

Structure of Solids

Crystalline: atoms, ions, or molecules are well ordered. Have a well-defined melting point. Often the solid has regular shapes.

Amorphous: no order to the particles. Examples are glass and rubber. Have no defined mp; they soften over a range of temperatures (important for glass blowing)

Unit cell: crystal lattice

In a brick wall there is a repeating pattern, as there is with most wallpaper

In a crystalline solid there is a repeating pattern - the unit cell. The unit cell repeats to make the crystal lattice

The seven unit cells

Three Cubic lattices

Close Packing

Another way of looking at it

Diamonds are a …..

Some ionic solids- lattice decided by size & charge

From sea to shining ...array of metal ions in a sea of electrons

A sea of valence electrons

Electrons not tightly held but can move

Explains electrical conduction

Also explains optical properties - most metals “shine”