1. Liquids, Solids, and Intermolecular Forces
Suggested HW: (Ch 15) 13-17

2. Important Recap
Consider the Lewis structures of the following molecules and identify them as polar or nonpolar, and justify your answer:
SO3
SO2
CH4
CH3F
CH3CH2CH3
NH3
CO2
TeCl4

3. States of Matter Differ By Intermolecular Distance
The state of a substance at a given temperature and pressure is determined by two factors:
Thermal energy of the molecules
Intermolecular forces (called Van der walls forces) between molecules

4. Intermolecular Forces
Intermolecular forces are the attractive forces that exist between molecules. These are not the same as intramolecular forces (i.e. bonds) which exist within a molecule.
Intermolecular forces are much weaker than intramolecular forces.
Consider water
Water would have to be heated to thousands of degrees to break the O-H covalent bonds.
However, the intermolecular forces can be overcome simply by boiling the water (100oC) which separates the molecules.
Nonetheless, intermolecular forces are very important and dictate many of the important physical characteristics of molecules

5. Intermolecular Forces: Coulombic Attractions
As you recall, ionic compounds are solids at room temperature. There are ion-ion attractions in ionic compounds.
The coulombic force that holds ions together is very strong. Coulombic attractions are the strongest of all intermolecular forces.
Therefore, all ionic compounds have very high melting/boiling points, and are solid at room temperature.

6. - - + + Dipole-Dipole Forces
Polar molecules attract one another. This type of intermolecular force is called dipole-dipole attraction. δ + δ - δ + δ -
Covalent bond: Very Strong
Dipole-dipole interaction:
Polar molecules will orient themselves in a way to maximize these attractions. The strength of these attractions increases with increasing polarity. Polar molecules have higher melting/boiling points than non polar ones.

7. Intermolecular Forces: Dipole-Dipole Forces
The magnitude of the melting/boiling temperatures of various substances reflect how strongly the molecules attract one another.
The more strongly the molecules attract, the harder it is to separate them. Hence, the higher the melting/boiling temperatures.
Recall polarity from chapter 8. Any molecule with a net dipole is polar. δ - + δ H Cl
Partial positive character
Partial negative character

8. Ion-Dipole Interactions
When salts are dissolved in water, the water dipoles are oriented around the ions.
This strong interaction between formal and partial charges is why adding salt to water raises its boiling temperature, or why salting snow causes it to melt.
Solution
Boiling Temp. (oC)
Boiling Temp. (oF)
Pure Water
100.0
212.0
Satd. NaCl (aq)
108.7
227.7

9. Hydrogen Bonding
There is a special, unusual, very strong type of dipole-dipole interaction known as hydrogen bonding
Hydrogen bonding is a dipole interaction that exists only between the H atom in an H—F, H—O, or H—N bond on one molecule and an adjacent lone pair on another F, O, or N atom in another molecule
Because hydrogen atoms are so small, the large partial positive charge induced on H when bonding with these highly electronegative elements is highly concentrated. Therefore, it strongly attracts neighboring N, O, and F atoms, and can approach them very closely.

10. δ-

11. Hydrogen Bonding Causes Abnormalities in Boiling Point Trend
Actual Boiling Temp
of Water (100oC)
Predicted Boiling Temp
of Water (-90oC)
Without hydrogen bonding, all water would be gas (steam) at normal room temperature!!!

12. Hydrogen Bonding
For example, consider ethanol and dimethyl ether (C2H6O, MW 46 g/mol)
The –OH bond in ethanol is susceptible to hydrogen bonding, whereas the C-O-C bond in dimethyl ether is not.
The strong attraction between these –OH groups results in the massive difference in boiling point between these substances.
ethanol
BP: 78.3oC
dimethyl ether
BP: oC

13. Structure and Density of Ice
Water is one of the few compounds that is less dense in its solid phase than its liquid phase.
This is due to hydrogen bonding.
In liquid water, 80% of the atoms are H-bonded. In ice, 100% are H-bonded.
To maximize H-bonding, the water molecules in ice spread out.
Therefore, we have the same mass of water, with a larger volume. Since ρ=(mass/volume), ρ decreases.

14. London Dispersion Forces
With nonpolar molecules, there are no dipoles, so we would not expect to see dipole-dipole interactions. Despite this, intermolecular interactions have still been observed.
For example, nonpolar gases like Helium can be liquified, but how can this happen? What force brings the He atoms together?
Fritz London, a physicist, proposed that the motion of electrons in a nonpolar molecule can create instantaneous dipoles

15. London Dispersion Forces
Lets take a Helium atom. At some moment in time, the electrons are spread out within the atom
However, because electrons are constantly moving, electrons can temporarily end up on the same side of the atom, creating an instantaneous dipole. e- 2+ δ- e- δ+ 2+ e-

16. London Dispersion Forces
These dipoles induce other dipoles on neighboring atoms, the cycle repeats indefinitely. This leads to the formation of a condensed phase. e- 2+
repel
attract e- 2+ e- δ- δ+ 2+ e- e- e- e- δ- δ+ δ- δ+ δ- δ+ 2+ 2+ 2+ e- e- e-

17. All substances have dispersion forces
London Dispersions
C5H12
C12H26
C18H38
The ease of the electron distortion is called polarizability. Bigger molecules are more polarizable, so they are more prone to instantaneous dipoles.
Hence, London dispersion forces increase with increasing molar mass because heavier atoms/molecules have more electrons, and are therefore are more polarizable.
All substances have dispersion forces

18. Boiling Points Increase With Increasing Strength of London Dispersion Forces

19. Constitutional Isomers and the Importance of Shape
309 K
282 K
301 K
Shown above are three molecules all having the formula C5H12, but different structures (constitutional isomers). Despite having the same molar masses, the boiling points are very different. Why?
Longer, flatter arrangement allows for more interaction of dipoles (stronger LD force)

20. Molecular Geometry of Fats And Oils
A typical nutritional label breaks down the types of fats present in a food to include saturated fats and unsaturated fats
A saturated fat is a long chain fatty acid containing no C=C double bonds
Unsaturated fats contain one (mono) or more (poly) double bonds
Chemically, fats and oils differ on this basis, although these terms are used interchangeably in nutrition. Fats are saturated fatty acids and are solid at room temperature. Unsaturated fats are technically oils, as they are liquids.
Oils can be converted into fats by breaking the C=C double bonds by a process called hydrogenation (ex. Margarine)

21. Molecular Geometry of Fats And Oils
Unsaturated fats can be separated into two classes, trans and cis. These differ by the orientation of terminal atoms about the C=C double bond.
trans
cis
The structure of trans oils allows the molecules to stack, whereas the cis structure does not. This stacking results in enhancement of the London dispersion force, which leads to the formation of a solid fat. This causes clogging of arteries.
trans stacking
cis stacking

22. Molecular Geometry of Fats And Oils
Saturated fats are considered to be worse for heart health than unsaturated fats because of the lack of “kinks” in the carbon chain, which makes it easier for the molecules to align and stack. n
clogged artery

23. Like Substances Mix
Solubility is also influenced by the strengths of intermolecular forces. To predict if a given solute will be soluble in a given solvent, we need to consider the strength and number of interactions.
If the solute-solute interactions are too strong, or if the solvent-solvent interactions are too strong, the molecules will not be able to separate to create a mixture.
Strong solute-solvent interactions are required for substances to mix

24. Like Substances Mix Polar substances are soluble in polar substances.
Nonpolar substances are soluble in nonpolar substances.
Polar and nonpolar substances DO NOT MIX (ex. oil and water)
E.G. Oil and water; Dipole-dipole and H-bond forces between water molecules are MUCH stronger than the London dispersion forces between oil and water
Water molecules would rather associate with other water molecules than to associate with oil.

25. Competing Forces: Molecules with both polar and nonpolar groups
The table below shows how increasing the size of the nonpolar (hydrophobic) portion of a ketone reduces its solubility in water. The additional hydrophobic groups strengthen solute-solute interaction through LD without contributing to the solute-solvent DD attractions of the polar (hydrophilic) portion

26. Competing Forces: Soaps
Soaps are a class of molecules called surfactants, molecules having long hydrophobic chains (tails) and bulky polar ends (heads)
The structure of the active ingredient of “Tide” is shown below. The tails bind to oil and dirt through LD, the polar heads solubilize the dirt through formation of a micelle (see left)

27. Examples Which of the following groups of compounds will mix?
O2 and CH3CH2CH2CH3 (butane)
SO2 and CCl4
H2O and CH3OH
K2S and H2O
H2O and CH3CH2NH2
Arrrange in order of increasing boiling point
SeO2, F2, Xe, NaCl, SeO3, CH2NH, CH3OH
Give two reasons why NH3 has such a lower BPo than H2O?

28. Intermolecular Forces At Work: Surface Tension
Imagine a body of liquid. The liquid molecules in the bulk have attractive intermolecular forces on all sides, pulling in all directions. These interactions minimize the energy of these molecules.
Molecules at the surface have less interactions, which makes them less stable due to their higher potential energy. The net force on the surface molecules is inward, and they cluster close together to minimize surface area. In other words, they want to minimize interactions with the air molecules above them.

29. Intermolecular Forces At Work: Surface Tension
The surface tension of a liquid is the energy required to increase its surface area by a unit amount (J/cm2 or mJ/m2). The surface tension of a liquid is what causes its surface to create a “skin” that resists penetration, or assume a droplet shape.

30. Intermolecular Forces At Work: Viscosity
Viscosity is a measure of a liquid’s resistance to flow and is a temperature dependent property (e.g. Maple syrup is more viscous than water). Stronger intermolecular forces lead to higher viscosity. Measured in the unit of centipoise (cP), where 1.00cP is the viscosity of water at 20oC.

31. Intermolecular Forces At Work: Viscosity

32. Intermolecular Forces At Work: Capillary Action
Closely related to surface tension, capillary action is the ability of a liquid to flow against gravity up a tube.
This is caused by a combination of intermolecular forces between neighboring liquid molecules (cohesive forces), as well as intermolecular forces between the liquid molecules and the molecules of the surface of the tube (adhesive forces).
If the adhesive forces are stronger than the cohesive forces, the liquid rises.

33. Intermolecular Forces At Work: Capillary Action
A consequence of capillary action is the formation of a meniscus (the shape formed by a liquid in a tube).
In systems with strong adhesion forces, the liquid molecules rise up the sides of the tube, forming a concave meniscus.
In systems with strong cohesion forces, the liquid molecules crowd around the interior to maximize their interactions, creating a convex meniscus.

34. Summary We can arrange the intermolecular forces by relative strength:
Coulombic attraction
Ion-dipole/Hydrogen Bonding
Dipole-Dipole
*London Dispersion
The list above is a general trend. For very large molecules, LD forces can exceed the strength of dipole-based forces.
ex. CH3CH2CH2CH2CH2CH2CH2OH) is nonpolar because the LD forces dominate.