1. Intermolecular forces
Liquids and Solids
Intermolecular forces

2. Chapter objectives
Physical properties of liquids: Vapor pressure and boiling
Intermolecular forces

3. Physical Property: Interactions Between Molecules
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

4. 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

5. Forces of Attraction within a Liquid
Cohesive Forces = forces that try to hold the liquid molecules to each other
Molecule  Molecule
surface tension
Adhesive Forces = forces that bind a substance to a surface
Molecule  Surface
capillary action
meniscus

6. Surface Tension
Surface tension: the tendency of liquids to minimize their surface. Cause: intermolecular force
liquids to have a surface that resists penetration
Paper clip (denser than water) can float on water

7. Viscosity
some liquids flow more easily than others: Soda more fluidy than Syrup
Viscosity : the resistance of a liquid to flow. Syrup is more viscous than Soda
 Attractive forces between the molecules (intermolecular forces)

8. Evaporation and Vapor Vapor: gaseous form from liquid
Evaporation : molecules of a liquid breaking free from the surface: Liquid  Gas
also known as vaporization
Physical change
Vapor: gaseous form from liquid

9. Evaporation: Liquid Molecule escape into Gas
Evaporation happens at the surface
molecules on the Surface experience a smaller net attractive force than molecules in the Interior
NOT all the surface molecules escape at once: only the ones with sufficient kinetic energy (fast enough) to overcome the attractions will escape

10. Condensation: Gas  Liquid
Condensation : the vapor molecules (gas state) may bump into and stick to the surface of the container or get recaptured by the liquid.
Physical change : Gas  Liquid
Condensation occurs for molecules with less kinetic energy and/or collides to surface

11. Evaporation vs. Condensation: Dynamic Equilibrium
Evaporation and Condensation are opposite processes
In a sealed container, eventually, the rate of evaporation and condensation in the container will be the same
Rate evaporation = Ratecondensation
Dynamic equilibrium : opposite processes that occur at the same rate in the same system
The amount of vapor vs. liquid appears constant

12. 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)

13. Vapor Pressure Pvap
once equilibrium is reached, then the amount of vapor (mole of vapor, nvap) in the container will remain the same
as long as you don’t change the conditions
Vapor pressure: the partial pressure exerted by the vapor of the liquid.
Depending on the temperature and strength of intermolecular attractions

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

15. Boiling and Boiling Point (b.p.)
Boiling: vapor pressure of the liquid is the same as the atmospheric pressure. Pvap = Pair
Rapid evaporation
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

16. Vapor pressure at given temperature vs. Normal Boiling point
At the same temperature, different liquids have different vapor pressure (volatility)
Volatile: Liquids having high vapor pressure
Liquids having higher vapor pressure will have lower normal boiling points

17. Energy flow: Evaporation vs. Condensation
Evaporation: Liquid _______ 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 cools your skin
Condensation: Gas _______ heat to its surroundings to reduce its temperature  The surroundings warms up
Exothermic: heat flows out of a system into the surroundings

18. Heat of Vaporization
Definition: the amount of heat needed to vaporize one mole of a liquid.
DHvap + Liquid  Gas
it requires 40.7 kJ of heat to vaporize one mole of water at 100°C
_____thermic
since Condensation (Gas  Liquid + DHvap ) is the opposite process to evaporation, ____thermic.
DHcond = -DHvap
DHvap is a Conversion Factor (unit: kJ/mol).

19. g H2O  mol H2O  heat 5.56 mole H2O 261 kJ
Example: Calculate the amount of heat is required to vaporize 100. g water at its boiling point. Molar heat of vaporization = kJ/mol
Given: 100. g water
Find: heat
CF: 40.7 kJ = 1 mol; g = 1 mol
g H2O
 mol H2O
 heat
5.56 mole H2O
261 kJ

20. Temperature and Melting
For solid, temperature increases until it reaches the melting point. Ice melts at 0°C.
During melting: the temperature remains _________ until it all turns to a liquid. solid  liquid
Why temperature remains constant? All the added heat is for overcoming the attractive forces in the solid, not increase the temperature

21. Energy of Melting and Freezing
Melting: Solid absorbs heat from its surroundings: _____thermic
as Heat flows out of the surroundings  the surroundings cool off
as ice in your drink melts, it cause the liquid to cool
Freezing: Liquid releases heat into its surroundings: _____thermic
as heat flows into the surroundings  the surroundings warm up

22. Heat of Fusion
Definition: heat needed to melt one mole of a solid DHfus
since freezing (crystallization) is the opposite process to melting, the same amount of energy transferred is the same, but in the opposite direction
DHcrystal = -Dhfusion
Heat of fusion (kJ/mol) can be used as conversion factor to calculate heat in melting/freezing problems

23. Heats of Fusion of Several Substances
Liquid
Chemical
Formula
Melting Point, °C
DHfusion, (kJ/mol)
water
H2O
0.00
6.02
isopropyl alcohol
C3H7OH
-89.5
5.37
acetone
C3H6O
-94.8
5.69
diethyl ether
C4H10O
-116.3
7.27

24. 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.

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

26. Intermolecular Forces (IMF) affects 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.
Stronger IMF also increases melting point

27. Why are molecules attracted to each other?
 Attractive forces between opposite electric charges
Ionic: + ion to – ion: Na+ Cl–
Molecular: (+) end of polar molecule to (-) end of polar molecule. HOd–-Hd+ Od–-Hd+
Ionic or Molecular: larger charge  stronger attraction:
Mg2+ Cl– > Na+ Cl–
Fd–-Hd+  Cld–-Hd+ > Cld–-Hd+  Cld–-Hd+
How about nonpolar molecules?

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

29. Dispersion Forces also: London Forces or Induced Dipoles
 Electrons on one molecule distorting the electron cloud on another
ALL molecules have Dispersion Forces
Dispersion force is especially important among nonpolar molecules + - + - + -

30. Dispersion Forces: Instantaneous Dipoles
Nonpolar
Somewhat polar
Polar

31. Strength of Dispersion Force
More electrons (more molar mass): Electrons can move more easily within a molecule: =O < =S -F < -Cl < -I
Example: CF4 (b.p. -127°C) vs. CBr4 (b.p. 190°C)
Larger molecules.
Example: CH4 (b.p. -161°C) vs. C2H6 (b.p. -89°C)
Why? more electrons + electron farther from the nuclei  the larger the dipole that can be induced

32. Practice: Which has higher boiling point?
Hint: Compare
molecular polarity (use VSEPR theory)
the number of electrons in each molecule/atom
CH4 vs. CCl4
F2 vs. Br2
Liquid argon vs. liquid xenon

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

34. 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

35. Attractive Forces Dispersion Forces – all molecules_ +
Dipole-to-Dipole Forces – polar molecules

36. Why iodine chloride (ICl, molar mass = 162, b. p
Why iodine chloride (ICl, molar mass = 162, b.p. 97 C) has higher boiling point than bromine (Br2, molar mass = 160, b.p. 59 C)?
Dispersion force:
ICl (70 e-) vs. Br2 (70 e-)
Dipole moments:
ICl (polar molec.) vs. Br2 (nonpolar molec.)
Hint: compare
the number of electrons in each molecule/atom
molecular polarity (use VSEPR theory)

37. Hydrogen Bonding  Hydrogen Bond
Molecules that have HF, -OH or -NH groups have particularly strong intermolecular attractions
unusually high melting and boiling points
unusually high solubility in water
Not for all molecules with hydrogen atom
 Hydrogen Bond

38. Intermolecular H-Bonding

39. Cause of 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–-

40. 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

41. Hydrogen Bond in DNA double helix

42. 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

43. Why dimethyl ether (CH3OCH3, b. p
Why dimethyl ether (CH3OCH3, b.p. 249 K) has much lower boiling point than ethanol (CH3CH2OH, b.p. 351 K)?
Dispersion force:
CH3OCH3 (20 e-) vs. CH3CH2OH (20 e-)
Dipole moments:
CH3OCH3 (1.30 D) < CH3CH2OH (1.63 D)
Hydrogen bonding:
CH3OCH3 (absent) vs. CH3CH2OH (present)
Hint: compare
the number of electrons in each molecule/atom
molecular polarity (use VSEPR theory)
Hydrogen bonding?

44. 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
ligroin (hexane), toluene, kerosene, CCl4
if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition

45. Solubility between two liquids: Immiscible Liquids
Pentane (C5H12) (C-H and C-C bond, nonpolar substance)
is mixed with water (O-H bond, polar)
the two liquids separate
 they are more
attracted to their own
kind of molecule than
to the other.

46. Types of Crystalline Solids

47. 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

48. Ionic Crystalline Solids
Ionic solids: composite units are formula units. NaCl
Na+  Cl– Na+  Cl–
Held together by Electrostatic forces between Cation+ and Anion–
cations and anions 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

49. 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

50. Types of Atomic Solids

51. 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

52. Types of Atomic Solids Nonbonding
Nonbonding Atomic Solid: held together by dispersion forces
Xe  Xe  Xe  Xe
Dispersion forces are relatively weak,
 very low melting points and DHfusion

53. 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

54. Metallic Bonding Model of metallic bonding explain:
luster, malleability, ductility, electrical and thermal conductivity  the mobility of the electrons in the solid
the strength of the metallic bond  Charge and Size of the cations – so the melting points and DHfusion of metals vary as well

55. 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: seawater as largest sample of solution.
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: mostly as liquid
expands as it freezes
most substances contract as they freeze
causes ice to be less dense than liquid water

56. Practice: Which of the following pairs of substances does the first one have higher boiling point?
water vs. hydrogen sulfide
sulfur dioxide vs. carbon dioxide
liquid oxygen vs. liquid nitrogen
CH3CH2OH vs. CH3OCH3
CH3CH2OH vs. CH3CH2SH
hydrogen chloride vs. hydrogen fluoride
methane vs. carbon tetrafluoride

57. Practice: Which of the following pairs of substances does the first one have higher boiling point?
water vs. hydrogen sulfide
sulfur dioxide vs. carbon dioxide
liquid oxygen vs. liquid nitrogen
CH3CH2OH vs. CH3OCH3
CH3CH2OH vs. CH3CH2SH
hydrogen chloride vs. hydrogen fluoride
methane vs. carbon tetrafluoride
H-bond
Dipole-dipole
Dispersion
H-bond
H-bond
H-bond
Dispersion

58. Practice: Which of the following pairs of substances is the first one more volatile?
water vs. hydrogen sulfide
sulfur dioxide vs. carbon dioxide
liquid oxygen vs. liquid nitrogen
CH3CH2OH vs. CH3OCH3
CH3CH2OH vs. CH3CH2SH
hydrogen chloride vs. hydrogen fluoride
methane vs. carbon tetrafluoride

59. Practice: Which is Which?
Rock candy (crystalline sugar)
Gold nugget
Copper(I) sulfide solid
Diamond
Atomic solid
Molecular solid
Ionic solid
Covalent solid