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3. What holds particles together in liquids and solids?
• Gas: widely spaced, rapid random mo3on, low
density
• Liquid: closer together, randomly arranged
• Solid: Closely packed, fixed posi3on, rigid, high
density
5. Intermolecular Forces
• Intermolecular forces: Forces between
molecules that cause them to aggregate and
form solids or liquids.
Intermolecular vs.
Intramolecular forces
(bonds)
“inter”: between molecules
“intra”: within a molecule
7. INTERmolecular Forces
1. Ion-Dipole force
2. Dipole-dipole force
3. Hydrogen bonding (a special type of dipole-
dipole force)
4. London dispersion force (induced dipole-
induced dipole force)
11. 3. Hydrogen Bonding
• A par3cularly strong type of dipole-dipole
force
• Occurs between polar molecules that have a H
bound to a highly electronega3ve atom ( -N-H,
-O-H, H-F)
• Examples:
O
H H
H F
N
H
H H
C
H
H
H
O
H
12. Hydrogen Bonding
The hydrogen bond occurs between:
1) lone pairs on electronega3ve atom of one molecule, &
2) H on electronega3ve atom of another molecule.
13. Hydrogen Bonding in Water
Water has two H atoms
and 2 lone pairs à forms
a vast network.
16. 4. London Dispersion Force (induced dipole-
induced dipole)
• London dispersion force occurs when
temporary, instantaneous dipole in one
molecule induces a similar dipole in a
neighboring molecule (temporary, random
rearrangement of charge).
Br2 At any given instant
17. London Dispersion Force
• Occurs between all molecules, but especially
important in nonpolar molecules b/c it’s the
only intermolecular force available to them.
• Examples
Ø I2, H2, CH4, CH3CH2CH3
Ø Noble gases like He, Ne, Ar in very cold temp
18. Strength of London Dispersion Forces
• London dispersion force is the weakest
intermolecular force.
• London dispersion force increases with:
1) Greater molar mass (due to larger, more
polarizable electron cloud)
2) Greater surface area of molecule
19. London Dispersion Force: Effect of Surface Area
Smaller surface area
Smaller total London
dispersion force
Larger surface area
Larger total London
dispersion force
21. Intermolecular Forces and Boiling Point
• Greater the intermolecular force between
the molecules of a substance, higher the
boiling point of that substance.
• i.e., More heat is required to break the
intermolecular forces to change liquid to gas.
22. Effect of Dipole-Dipole Force on Boiling Point
NO dipole-dipole force Has dipole-dipole force
23. Effect of Hydrogen Bonding on Boiling Point
• Hydrogen bonding
makes the boiling
point of water very
high, rela3ve to
hydrides of other
group 6 elements.
Ex Probs
25. Phase Changes
• Phase change: Change in the physical state of
ma_er (gas, liquid, solid)
26. Phase Changes
• Mel3ng (Fusion): solid to liquid
• Freezing: liquid to solid
• Vaporiza3on: liquid to gas
• Condensa3on: gas to liquid
• Sublima3on: solid to gas
• Deposi3on: gas to solid
27. Heat in Phase Changes
• Energy is transferred as heat during phase
changes.
• Heat (q) absorbed or released by a process* =
Enthalpy change (ΔH) = Hfinal – Hini3al
• ΔH = q (*under constant pressure)
28. Exothermic vs. Endothermic Process
• Exothermic process (-ΔH): a process that releases
heat
• Endothermic process (+ΔH): a process that
absorbs heat
29. Enthalpy Change of Phase Changes
• Mel3ng (Fusion): solid to liquid à endothermic +ΔH
• Freezing: liquid to solid à exothermic -ΔH
• Vaporiza3on: liquid to gas à endothermic +ΔH
• Condensa3on: gas to liquid à exothermic -ΔH
• Sublima3on: solid to gas à endothermic +ΔH
• Deposi3on: gas to solid à exothermic -ΔH
30. Enthalpy Changes of Phase Changes
• Phase changes from Solid à Liquid à Gas
are endothermic because heat is needed to
break the intermolecular forces.
33. Temperature of Phase Change
• Mel3ng point (mp): the temperature at which
solid turns into liquid (temperature of
transi3on between solid and liquid) = freezing
point
• Boiling point (bp): the temperature at which
liquid turns into gas (the temperature of
transi3on between liquid and gas)
34. Heating Curve for Water
• Within a single phase: temperature increases
when heat is added, b/c…
• the heat increases the kine3c energy (energy
of mo3on) of the par3cles.
KEavg ∝ T
• Can calculate the temperature change through
q = mCΔT.
35. Heating Curve for Water
• During a phase change: temperature remains
constant 3ll phase change is completed for the
whole sample as heat is added, because…
the added heat is breaking apart the
intermolecular forces (a_rac3ve forces
between par3cles).
36. Heats of Phase Changes
• Heat of fusion (ΔHfus): quan3ty of heat required to
completely melt one mole of a substance once it has
reached its mel3ng point
q = moles x ΔHfus
• Heat of vaporiza3on (ΔHvap): quan3ty of heat
required to completely vaporize one mole of a
liquid once it has reached its boiling point.
q = moles x ΔHvap
38. Vapor Pressure of a Liquid
Vapor: the gaseous state of a substance that is
normally liquid (or solid) at room temperature
Vapor pressure: the pressure exerted by a vapor in
equilibrium with its liquid phase in a closed system
During
evapora3on,
molecules
escape from
surface of
liquid.
39. Boiling Point
• When temperature reaches boiling point,
molecules in interior of liquid have enough KE
to escape as gas (bubbles).
• Boiling point: the temperature at which
vapor pressure = atmospheric pressure
Bubbles can form and
rise since the vapor
pressure can overcome
the atmospheric
pressure.
40. Water: A Unique Liquid
Water is a polar molecule that has strong
hydrogen bonding, which causes water to:
§ Have highest specific heat capacity of any liquid
§ Have unusually high heat of vaporiza3on (540
cal/g)
§ Be more dense as a liquid than as a solid
(unique)
41. Why Ice Floats
• In ice, water molecules are locked into posi3on
by hydrogen bonding à more open structure
than liquid water à less dense than liquid
water.
• Explains why ice floats, pipes burst when they
freeze
43. Structures of Solids
Two Major Categories of Solid Structures
1. Crystalline: Atoms, ions, or molecules are
ordered in well-defined 3D arrangements
2. Amorphous: Par3cles have no orderly structure.
