1. Intermolecular forces
Liquids and Solids

2. Chapter objectives
Understand the three intermolecular forces in pure liquid in relation to molecular structure/polarity
Understand the physical properties of liquids that is relevant to intermolecular force: vapor pressure and boiling

3. How Gecko can climb upside down on the ceiling?
Gecko are capable of climbing vertical or even upsidedown, utilizing micrometer size hairs to adhere to surfaces: Intermolecular force.
New research showed gecko can turn on and off the “stickiness”.

4. Intermolecular forces affect physical properties of solid & liquid
stronger intermolecular forces increases surface tension and viscosity
stronger intermolecular forces reduces vapor pressure (retaining molecules in liquid state), thus increases boiling point
likewise with melting point

5. Surface Tension
Surface tension: Tendency of liquid to minimize total surface.
Cause: Molecules on the surface attract to each other.
Paperclip (made of iron, density >> water) should sink in water.

6. Viscosity Viscosity: Internal friction within the liquid.
Stronger intermolecular forces increases viscosity
Viscous liquid takes longer to flow.
Honey has higher viscosity than water
Used motor oil has less viscosity

7. Intermolecular forces affect physical properties of solid and liquid
Stronger intermolecular force in the liquid prevent liquid molecules from escaping into gas state
 Molecules need more heat energy to escape from the liquid (higher temperature)
 Leading to ________ (higher, lower) boiling point.

8. Intermolecular forces in Pure liquids
Dipole-dipole force
Dispersion force (aka London force)
Hydrogen bonding

9. Permanent Dipoles Chapter 4:
Electronegativity difference & Molecular Geometry  some molecules have a Permanent Dipole: (+)  (-)
all polar molecules have a permanent dipole. H2O, NH3, HCl, etc.

10. Dipole-to-Dipole Attraction
Polar molecules have a permanent dipole
a + end and a – end
the + end of one molecule will be attracted to the – end of another
Similar to attraction between two separated bar magnets

11. Polar Molecules Can be attracted by Charge
Demo:
Note: Only works for polar liquid. Nonpolar liquid, such as gasoline, doesn’t work.
When two polar molecules are coming close, there are both attraction and repulsion between them.
Since the Coulombic force depends on the distance between the charge, the attraction between two polar molecules is STRONGER than repulsion, leading to intermolecular attractive force.

12. Dispersion Forces
Nonpolar molecules also attract each other: London Forces or Induced Dipoles
Cause: Electrons on one molecule distorting the electron cloud on another
ALL molecules have Dispersion Forces
Dispersion force is especially important among nonpolar molecules + - + - + -

13. Dispersion Forces: Instantaneous Dipoles
Nonpolar
Somewhat polar
Polar

14. Dispersion Force: Strength
Strength of the dispersion force gets Larger with larger molecules
Electron mobility: how easily the electrons can move within a molecule, or be polarized. =O < =S -F < -Cl < -I
more electrons + electron farther from the nuclei  the larger the dipole that can be induced

15. Dispersion Force: Nonpolar molecules F2, Cl2, Br2, I2
Larger molar mass, stronger intermolecular forces leads to higher boiling point, higher melting point.

16. Attractive Forces: Comparison
Dispersion Forces – all molecules _ +
Dipole-to-Dipole Forces – polar molecules

17. Beyond Dipole-Dipole and London Dispersion forces?
London force: high molar mass, stronger force, higher boiling point?
Anomaly for water and HF

18. Hydrogen Bonding
Molecules that have HF, -OH or -NH groups have particularly strong intermolecular attractions
unusually high melting and boiling points
unusually high solubility in water
 Hydrogen Bond

19. Intermolecular H-Bonding

20. Hydrogen Bonding
A very electronegative atom X (X = F, O, N) is bonded to hydrogen, the bonding electrons is pulled toward X.
Xd–-Hd+
Since hydrogen has no other electrons, the nucleus becomes deshielded (“stripped”): -Hd+
exposing the proton
The exposed proton Hd+ (center of positive charge) attracting all the electron clouds from neighboring molecules Xd–-Hd+  Yd–-

21. H-Bonds vs. Chemical Bonds
Hydrogen bonds are not chemical bonds
Hydrogen bonds are attractive forces between molecules
Chemical bonds are attractive forces that make molecules

22. Hydrogen Bonding in DNA double helix

23. Types of Intermolecular Forces
Type of Force
Relative Strength
Present in
Example
DispersionForce
weak, but increases with molar mass
all atoms and molecules H2
Dipole – Dipole Force
moderate
only polar molecules
HCl
Hydrogen Bond
strong
molecules having H bonded to F, O or N HF

24. Attractive Forces and Solubility
Like dissolves Like
miscible = liquids that do not separate
Polar molecules dissolve in Polar solvents
water, alcohol, isopropanol, CH2Cl2
H-bond: molecules with O or N higher solubility in H2O
Nonpolar molecules dissolve in nonpolar solvents
Gasoline, Paint thinner, toluene, kerosene, CCl4
if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition

25. Solubility between two liquids: Miscible vs. Immiscible
Water and alcohol can mix at any ratio
Immiscible: Pentane (C5H12) (C-H and C-C bond, nonpolar)
is mixed with water (O-H bond, polar)
the two liquids separate

26. Interaction Between Molecules Affects Physical Property
Many of the phenomena we observe are related to interactions between molecules that do not involve a chemical reaction
your taste and smell organs work because molecules interact with the receptor molecule sites in your tongue and nose
The volatility of liquids depends on the IMF.

27. Structure Determines Properties: Solids, Liquid and Gases

28. Why is Sugar a Solid But Water is a Liquid?
The state a material exists in depends on the attraction between molecules and their ability to overcome the attraction
The attractive forces between Ions or Molecules  Their structure
the attractions are electrostatic
depend on shape, polarity, etc.
The ability of the molecules to overcome the attraction  Kinetic energy they possess

29. Escaping from the Surface
Evaporation : molecules of a liquid breaking free from the surface: Liquid  Gas
also known as vaporization
Physical change
a substance is converted from its liquid form to its gaseous form
the gaseous form is called a vapor

30. Evaporation: Liquid  Gas
Molecules of the liquid mix with and dissolve in the air
happens at the surface
molecules on the Surface experience a smaller net attractive force than molecules in the Interior
but all the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape

31. Condensation: Gas  Liquid
in a closed container, after a liquid evaporates, the vapor molecules are trapped and may eventually turn into liquid
Condensation : the vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid.
Physical change : Gas  Liquid

32. Dynamic Equilibrium
Evaporation and Condensation are opposite processes
eventually, the rate of evaporation and condensation in the container will be the same
Dynamic equilibrium : opposite processes that occur at the same rate in the same system

33. Evaporation  Condensation
Water is just added to the flask and it is capped, all the water molecules are in the liquid.
Eventually,
Rateevap = Ratecondsn
The air in the flask
is now saturated
with water vapor.
Shortly, the water
starts to evaporate.
Speed of evaporation >> Speed of condensation
(Rateevap >> Ratecondsn)

34. Vapor Pressure Pvap
once equilibrium is reached, then the amount of vapor (mole nvap) in the container will remain the same
Vapor pressure: Pressure exerted by the vapor of the liquid when equilibrium is reached between liquid and gas states.
Depending on the temperature and strength of intermolecular forces (IMF): Stronger IMF, more ____________ for liquid to become vapor, ___________ (higher, lower) Pvap.

35. Vapor Pressure increases as temperature increases
ethanol
ether
normal boiling point
water

36. Boiling and Boiling Point (b.p.)
Boiling: vapor pressure of the liquid is the same as the atmospheric pressure. Liquid  Gas. Pvap = Pair
Boiling point: the temperature for boiling process
normal boiling point: temperature when Pair = 1 atm b.p. of water is 100°C
b.p. depends on Pair
the temperature of boiling water on the top of a mountain will be cooler than boiling water at sea level
On top of Mount Whitney, b.p. of water is about 84°C

37. Vapor pressure at given temperature vs. Normal Boiling point
At the same temperature, different liquids have different vapor pressure (volatility)
Liquids having higher vapor pressure are normally called “more volatile”
Liquids having higher vapor pressure will have lower normal boiling points

38. Energy flow: Evaporation vs. Condensation
Evaporation: Liquid absorbs heat from its surroundings to evaporate  The surroundings cool off
Endothermic: heat flows into a system from the surroundings
as alcohol evaporates off your skin, it causes your skin to cool
Condensation: Gas releases heat to its surroundings to reduce its temperature  The surroundings warms up
Exothermic: heat flows out of a system into the surroundings

39. Temperature and Melting
For solid, temperature increases until it reaches the melting point. Ice melts at 0°C.
During melting: the temperature remains the same until it all turns to a liquid. solid  liquid
all the Energy from the heat source is for overcoming the attractive forces in the solid, not increase the temperature

40. Heating Curve: phase changes during heating solid ice at 1 atm
s+l s g l
l+g

41. Sublimation vs. Deposition
Sublimation: the Solid form changes directly to the Gaseous form. Solid  Gas
without going through the liquid form
Dry ice (solid CO2)  gas CO2
like melting, sublimation is endothermic
Deposition is the reverse of Sublimation, exothermic.

42. Types of Crystalline Solids

43. Molecular Crystalline Solids
Molecular solid: composite units are molecules.
CO2  CO2 
H2O  H2O  H2O
Held together by intermolecular attractive forces
dispersion, dipole-dipole, or H-bonding
generally low melting points and DHfusion

44. Ionic Crystalline Solids
Ionic solids: composite units are formula units. NaCl
Na+  Cl– Na+  Cl–
Held together by Electrostatic forces between Cation+ and Anion–
arranged in a geometric pattern called a crystal lattice to maximize attractions
generally higher melting points and DHfusion than molecular solids
because ionic bonds are stronger than intermolecular forces

45. Atomic Crystalline Solids
Atomic solids: composite units are individual atoms
Xe  Xe  Xe  Xe
Held together by either covalent bonds, dispersion forces or metallic bonds
melting points and DHfusion vary depending on the attractive forces between the atoms

46. Types of Atomic Solids

47. Types of Atomic Solids Covalent
Covalent Atomic Solids : atoms attached by covalent bonds. Diamond Carbon (tetrahedral, C-C bond).
effectively, the entire solid is one, giant molecule
Covalent bonds are strong
 very High melting points and DHfusion
 High hardness

48. Types of Atomic Solids Nonbonding
Nonbonding Atomic Solid: held together by dispersion forces. Xenon solid (at low temperature)
Xe  Xe  Xe  Xe
Dispersion forces are relatively weak,
 very low melting points and DHfusion

49. Types of Atomic Solids Metallic
Metallic solids: held together by metallic bonds
How: metal atoms release some of their electrons to be shared by all the other atoms in the crystal
Metallic bond: the attraction of the metal Cations M+ for the mobile electrons e-
often described as “islands of cations in a sea of electrons”

50. Water: A Unique and Important Substance
found in all 3 states on the Earth: Ice, Liquid, Vapor
the most common solvent (liquid) found in nature
without water, life as we know it could not exist
the search for extraterrestrial life starts with the search for water
relatively high boiling point
expands as it freezes
most substances contract as they freeze
causes ice to be less dense than liquid water