1. Intermolecular forces
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Questions
Why do some solids dissolve in water but others do not?
Why are some substances gases at room temperature, but others are liquid or solid?
What gives metals the ability to conduct electricity, what makes non-metals brittle?
The answers have to do with …
Intermolecular forces

2. Intermolecular forces
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Intermolecular forces
Overview
There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces
We have already looked at intramolecular bonds (ionic, polar, non-polar)
Intermolecular forces (IMF) have to do with the attraction between molecules (vs. the attraction between atoms in a molecule)

3. Ionic, Dipole - Dipole attractions
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We have seen that molecules can have a separation of charge
This happens in both ionic and polar bonds (the greater the EN, the greater the dipoles) H Cl + –
Molecules are attracted to each other in a compound by these +σ and - σ forces + – + – + – + –

4. H - bonding
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H-bonding is a special type of dipole - dipole attraction that is very strong
It occurs when N, O, or F are bonded to H
Q- Calculate the EN for HCl and H2O
A- HCl = = 0.8, H2O = = 1.4
The high EN of NH, OH, and HF bonds cause these to be strong forces (about 5x stronger than normal dipole-dipole forces)
They are given a special name (H-bonding) because compounds containing these bonds are important in biological systems

5. Hydrogen Bond Strongest of all “weak” forces
Is caused when H is bonded to F, O, or N
These are so electronegative that the H is a “naked nucleus” or bare proton
Very attractive!
Will bond to nearby electron pairs

6. Importance of Hydrogen Bonding
Biological systems – DNA, proteins
Water chemistry (MP, BP, specific heat, surface tension)
Density of ice

7. Density Most solids are more dense than liquid
Water is less dense because of hydrogen bonding
At 4°C, water becomes less dense
Important for life in winter
Causes lake turnover
Alum example

8. London forces
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Non-polar molecules do not have dipoles like polar molecules. How, then, can non-polar compounds form solids or liquids?
London forces are named after Fritz London (also called van der Waal forces)
London forces are due to small dipoles that exist in non-polar molecules
Because electrons are moving around in atoms there will be instants when the charge around an atom is not symmetrical
The resulting tiny dipoles cause attractions between atoms/molecules

9. London forces Instantaneous dipole: Induced dipole:
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Instantaneous dipole:
Induced dipole:
Eventually electrons are situated so that tiny dipoles form
A dipole forms in one atom or molecule, inducing a dipole in the other

10. London Dispersion Forces
All molecules have this
Only attraction in nonpolar molecules
How can Iodine be a solid?
Temporary lopsided charge builds up from random motion of electrons
Increases with mass – we say it has greater polarizability
Straight molecule is more polarizable than a curled up molecule – why?
Halogen Family is a great essay

11. Ionic, H-bonding, Dipole, or London?
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Ionic, H-bonding, Dipole, or London?
Details
Bond
Molecule
IMF
EN =
nonpolar
London
EN =
polar
dipole-dipole*
EN =
ionic
Ionic
ionic*
H + N,O,F
H-bonding*
Symmetrical molecule (any EN) --
*Since all compounds have London forces. London forces are also present. However, their affect is minor and overshadowed by the stronger forces present. Note: the term “polar” is used interchangeable with “polar covalent”. Likewise, “nonpolar” and “nonpolar covalent” mean the same thing.

12. Mixing oil and water
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Lets take a look at why oil and water don’t mix (oil is non-polar, water is polar) + – + –
The dipoles of water attract, pushing the oil (with no partial charge) out of the way: attractions win out over the tendency toward randomness + –

13. Properties of Liquids Viscosity “Slower than…..
Resistance of a liquid to flow
Time it as it goes through a small tube with gravity acting upon it.
Poise – 1g/cm-s
Trends – same substance – decreases with increasing temperature
series (same structure) – increases with increasing mass

14. Surface Tension How many drops on a penny? Uneven forces at surface
Acts like pond scum
Definition – energy needed to increase the surface area of a liquid by a certain amount
Water is high – why?
Called “cohesive” force – together
Water moving up a stem – adhesive force
Capillary acion – rise up a thin tube
Meniscus!

15. Phase Changes Solid to Liquid is called Heat of Fusion Hfus
For water, 6 kJ/mol
Liquid to Gas is called Heat of Vaporization
Hvap
For water, 40.7 kJ/mol
Hsub is sum of each

16. Heating Curve Try a problem
Remember - flat during phase change, temperature change when heating a single phase
Cooling is opposite

17. Supercooling
Happens with some liquids - remove heat and it doesn’t freeze when it should
Very unstable
May happen during hibernation

18. Critical Temperature
Highest temperature at which a liquid can form from a gas when pressure is applied.
Above this, the substance is called a supercritical fluid.
Gas just becomes more compressed.
Critical pressure - pressure at the critical temperature

19. Vapor pressure
Vapor pressure forms above any liquid if container is closed – why?
Equilibrium is reached
This is vapor pressure
Higher if forces holding liquid together are weak - called a volatile (fleeing) liquid

20. Boiling Point Temperature at which the VP equals atmospheric pressure
Normal BP - boiling point at 1 atm
Everest? Autoclave?

21. Phase Diagram Handout Look at lines Look at slope of AB
Freeze-drying - library book example

22. Water vs. CO2

23. Structure of Solids
Amorphous (rubber, plastics) - large or mixtures - no true structure
Crystalline - highly ordered structure
Crystalline solids have true melting points

24. Unit Cell Repeating unit of a solid 7 types – (6-sided parallelograms)
Ni, Na, NaCl
Array of points in the crystal lattice

25. 3 cubic unit cells

26. Total Atoms for each unit cell

27. Packing
Spheres naturally pack hexagonally
Animation

28. Bonding Shown by x-ray diffraction Molecular - low MP
If unit packs well, mp can be high
Covalent Network Solid - very strong
Many covalent bonds in 3-D
Diamond, graphite, SiO2, SiC, BN
Ionic - greater charge, greater MP
Metallic solids - hexagonal close packed, mp varies

29. Network solids (covalent crystals)
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There are some compounds that do not have molecules, but instead are long chains of covalent bonds (E.g. diamond) C C C C C C C C C C C C C C C C C
This happens in 3 dimensions, creating a crystal
Because there are only covalent bonds, network solids are extraordinarily strong

30. Metallic crystals
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Metals normally occur as solids (high melting points).
Thus, there must be strong bonds between the atoms of metals causing them to bond
Bonding in metals and alloys is different from in other compounds: positive nuclei exist in a sea of electrons (this explains why metals conduct electricity) + –

31. Crystal types There are 6 types of intermolecular forces
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There are 6 types of intermolecular forces
These forces are associated with certain crystal types. By comparing solids we have a common frame of reference.
The crystal types and their basic units are
1) Network (covalently bonded atoms)
2) ionic (electrostatic attraction of ions),
3) metallic (positive nuclei in electron sea),
4) Molecular (electrostatic attraction of dipoles in molecules)
a) Polar (dipole-dipole and H-bonding)
b) Non-polar (London forces)

32. Properties of crystals
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Boiling and melting occur when the forces between molecules are overcome and a change of state occurs
The higher the force of attraction between molecules (IMF) the higher the melting/-boiling point (see previous slide for order)
Only metallic crystals conduct electricity in solid state (they also conduct in liquid state)
Ionic crystals will conduct electricity in molten state or dissolved because ions are free to move to positive and negative poles

33. Solubility of crystal types
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Solute = what is dissolving (e.g. salt)
Solvent = what it is dissolving in (e.g. water)
Strong attractions between the basic units of covalent crystals cause them to be insoluble.
Metallic crystals are likewise insoluble
The solubility of other crystals depends on solute and solvent characteristics
We will see that polar/ionic solutes dissolve in polar/ionic solvents and non-polar solutes dissolve in non-polar solvents
This is known as the like-dissolves-like rule

34. Crystals worksheet Place numbers in boxes according to descriptions
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Place numbers in boxes according to descriptions
Use pg and class notes, and internet as references
At end rows should add up to 126 each

35. Crystals worksheet answers
12/10/99 P A M E T C
Metallic 2 7 13 16 23 30 35
Network 3 10 11 20 21 29 32
Ionic 5 8 12 17 24 27 33
Molecular polar 4 9 14 18 22 28 31
Molecular non-polar 1 6 15 19 25 26 34
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36. Testing concepts
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Which attractions are stronger: intermolecular or intramolecular?
How many times stronger is a covalent bond compared to a dipole-dipole attraction?
What evidence is there that nonpolar molecules attract each other?
Suggest some ways that the dipoles in London forces are different from the dipoles in dipole-dipole attractions.
A) Which would have a lower boiling point: O2 or F2? Explain. B) Which would have a lower boiling point: NO or O2? Explain.

37. a) NH3, b) SF6, c) PCl3, d) LiCl, e) HBr, f) CO2
Which would you expect to have the higher melting point (or boiling point): C8H18 or C4H10? Explain.
What two factors causes hydrogen bonds to be so much stronger than typical dipole-dipole bonds?
So far we have discussed 4 kinds of intermolecular forces: ionic, dipole-dipole, hydrogen bonding, and London forces. What kind(s) of intermolecular forces are present in the following substances:
a) NH3, b) SF6, c) PCl3, d) LiCl, e) HBr, f) CO2
(hint: consider EN and molecular shape/polarity)
Challenge: Ethanol (CH3CH2OH) and dimethyl ether (CH3OCH3) have the same formula (C2H6O). Ethanol boils at 78 C, whereas dimethyl ether boils at -24 C. Explain why the boiling point of the ether is so much lower than the boiling point of ethanol.
Challenge: try answering the question on the next slide.
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38. Testing concepts Intramolecular are stronger.
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Testing concepts
Intramolecular are stronger.
A covalent bond is 100x stronger.
The molecules gather together as liquids or solids at low temperatures.
London forces
Are present in all compounds
Can occur between atoms or molecules
Are due to electron movement not to EN
Are transient in nature (dipole-dipole are more permanent).
London forces are weaker

39. 12/10/99
Testing concepts
A) F2 would be lower because it is smaller. Larger atoms/molecules can have their electron clouds more easily deformed and thus have stronger London attractions and higher melting/boiling points.
B) O2 because it has only London forces. NO has a small EN, giving it small dipoles.
C8H18 would have the higher melting/boiling point. This is a result of the many more sites available for London forces to form.
1) a large EN, 2) the small sizes of atoms.

40. Testing concepts a) NH3: Hydrogen bonding (H + N), London.
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Testing concepts
a) NH3: Hydrogen bonding (H + N), London.
b) SF6: London only (it is symmetrical).
c) PCl3: EN= Dipole-dipole, London.
d) LiCl: EN= Ionic, (London).
e) HBr: EN= Dipole-dipole, London.
f) CO2: London only (it is symmetrical)
Challenge: In ethanol, H and O are bonded (the large EN results in H-bonding). In dimethyl ether the O is bonded to C (a smaller EN results in a dipole-dipole attraction rather than hydrogen bonding).

41. H – bonding and boiling point
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why does BP as period , why are some BP high at period 2?

42. For more lessons, visit www.chalkbored.com
Testing concepts
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Boiling points increase down a group (as period increases) for two reasons: 1) EN tends to increase and 2) size increases. A larger size means greater London forces.
Boiling points are very high for H2O, HF, and NH3 because these are hydrogen bonds (high EN), creating large intermolecular forces
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