chemistry 100 – chapter 11 intermolecular forces
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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”
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