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Ch 11: Intermolecular Forces, Liquids, and Solids

Brown, LeMay AP Chemistry

2 11.1: Intermolecular Forces (IMF)   Intermolecular forces (IMF) are the forces that exist

between the molecules of a substance

  IMF are weaker than intramolecular forces (forces within molecules like covalent, metallic, ionic bonds)

  Inter means

  Intra means

3 11.1: Intermolecular Forces (IMF)

  Think about the states of matter. 1.  With your table discuss the strength of the

intermolecular forces in a state of matter. Rank the strength from weakest to strongest.

2.  Once you have your ranking, Use your understanding of the different states of matter to support why your ranking is correct.

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4 11.1: Intermolecular Forces (IMF) IMF Strength and States of Matter

Gases < Liquids < Solids

5 11.1: Intermolecular Forces (IMF)

Boiling Point & Melting Point and IMF strength:

Do Intermolecular Forces impact boiling and melting points of a substance? Explain your thoughts.

6 11.2: Types of IMF

Access the class website AP Chem Unit 5 Watch, read, and take notes on the first two

learning objects: 1. Intermolecular Forces 2. Hydrogen Bonding

*Draw the images from the simulations into your notes

I encourage you to discuss what you are seeing with your groups.

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7 11.2: Types of IMF What types of Intermolecular forces are found

in the following image? (Not only name the type but also the components involved in that interaction)

8 11.2: Types of IMF 1.  Electrostatic forces: act over larger distances in

accordance with Coulomb’s law

a.  Ion-ion forces (Ionic Bonding): strongest; found in ionic crystals (i.e. lattice energy)

9 b.  Ion-dipole: between an ion and a dipole (the partial charge on the end of a polar molecule)   Increase with increasing polarity of molecule

and increasing ion charge.

δ- δ+

δ+

δ- δ+

δ+ δ+

δ- δ+

Cl- δ- δ+

δ+

δ- δ+

δ+ δ+

δ- δ+

S2- <

Ex: Compare IMF in Cl- (aq) and S2- (aq).

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c.  Dipole-dipole: weakest electrostatic force; exist between neutral polar molecules

  Increase with increasing polarity (dipole moment) of molecule

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d.  Hydrogen bonds (or H-bonds):   H is unique among the elements because it

has a single e- that is also a valence e-. –  When this e- is “hogged” by a highly EN atom

(a very polar covalent bond), the H nucleus is partially exposed and becomes attracted to an e--rich atom nearby.

–  Generally stronger than dipole-dipole forces and dispersion forces (Will be talked about in a minute)

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  H-bonds form with H-X and X‘ from another molecule, where X and X' have high Electronegativity and X' possesses a lone pair of e-

  X = F, O, N (since most electronegative elements) on two molecules:

F-H

O-H

N-H

:F

:O

:N

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13   H-bonds between Amino Acids determine protein structure and strength

  Hold DNA strands together in double-helix

Nucleotide pairs form H-bonds DNA double helix

14   H-bonds explain why ice is less dense than water.

15 Ex: Boiling points of nonmetal hydrides

Boilin

g Po

ints

(ºC

)

Conclusions:

 Polar molecules have higher BP than nonpolar molecules

Therefore polar molecules have stronger IMF

 BP increases with increasing MW

Therefore heavier molecules have stronger IMF

 NH3, H2O, and HF have unusually high BP. Therefore H-bonds are stronger than dipole-dipole IMF

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16 2.  Inductive forces:   Arise from distortion of the e- cloud induced

by the electrical field produced by another particle or molecule nearby.

  London dispersion: between polar or nonpolar molecules or atoms –  Proposed by Fritz London in 1930 –  Must exist because nonpolar molecules form

liquids Fritz London(1900-1954)

17 How they form: 1.  Motion of e- creates an instantaneous dipole

moment, making it “temporarily polar”.

2.  Instantaneous dipole moment induces a dipole in an adjacent atom * Persist for about 10-14 or 10-15 second

Ex: two He atoms

18 Polarizability •  Ease with which the electron distribution in a

molecule is distorted •  Determines the strength of Dispersion

Forces •  Greater the polarizability of a molecule, the

more easily its electron cloud can be distorted to create a momentary dipole

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19 London dispersion forces increase with:   Increasing MW, # of e-, and # of atoms (increasing

# of e- orbitals to be distorted) Boiling points:

Effect of MW: Effect of # atoms: C5H12 36ºC Ne –246°C C6H14 69ºC CH4   –162°C C7H16 98ºC

  “Longer” shapes (more likely to interact with other molecules)

C5H12 isomers: n-pentane (straight chain) 309.4K neopentane (compacted) 282.7K

20 London Dispersion Forces •  Exists between all molecules

– Polar molecules experience dipole-dipole interactions along with dispersion forces (dipserions in polar molecules commonly contribute more to intermolecular attraction)

•  Comparing relative strengths of IMF in two substances: –  IF MW are similar dispersion forces are

similar. Dipole-dipole attractions overrule. (more polar stronger)

–  IF MW are greatly different dipersion forces determine strength of attractions. (Higher MW stronger)

Summary of IMF

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22 Ex: Identify all IMF present in a pure sample of each substance, then explain the boiling points.

BP(°C) IMF Explanation

HF 20

HCl -85

HBr -67

HI -35

Lowest MW/weakest London, but most polar/strongest

dipole-dipole and has H-bonds

Low MW/weak London, moderate polarity/dipole-dipole

and no H-bonds

Medium MW/medium London, moderate polarity/dipole-dipole

and no H-bonds

Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no

H-bonds

London, dipole-dipole, H-bonds

London, dipole-dipole

London, dipole-dipole

London, dipole-dipole

23 Practice Problem •  List the following substances in order of increasing boiling

points: BaCl2, H2, CO, HF, and Ne

24 11.3: Properties resulting from IMF 1.  Viscosity: resistance of a liquid to flow; the ease with

which individual molecules can move with respect to one another

2.  Surface tension: energy required to increase the surface area of a liquid

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3. Cohesion: attraction of molecules for other molecules of the same compound

4. Adhesion: attraction of molecules for a surface

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5.  Meniscus: curved upper surface of a liquid in a container; a relative measure of adhesive and cohesive forces Ex:

Hg H2O (cohesion rules) (adhesion rules)

27 11.4: Phase Changes Processes:   Endothermic:

melting (sl), vaporization (l g), sublimation (s g)

  Exothermic: condensation (g l), freezing (l s), deposition (gs)

I2 (s) and (g)

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29 Water: Enthalpy Diagram or Heating Curve

30 11.5: Vapor pressure

  A liquid will boil when the vapor pressure equals the atmospheric pressure, at any T above the triple point.

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31 11.6: Phase diagrams:

  Lines: 2 phases exist in equilibrium

  Triple point: all 3 phases exist together in equilibrium (T on graph)

  Critical point, or critical temperature & pressure: highest T and P at which a liquid can exist (C on graph)

32 Phase diagrams: H2O and CO2

  For most substances, increasing P will cause a gas to condense (or deposit), a liquid to freeze, and a solid to become more dense (to a limit.)

•  For H2O, increasing P will cause ice to melt. (Hydrogen Bonding)

33 11.7-8: Structures of solids   Amorphous: without orderly structure

Ex: rubber, glass

  Crystalline: repeating structure; have many different stacking patterns based on chemical formula, atomic or ionic sizes, and bonding

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Types of crystalline solids (Table 11.7)

Type Particles Forces Notable properties Examples

Atomic Atoms London dispersion

 Poor conductors  Very low MP

Ar (s),Kr (s)

Molecular Molecules (polar or

non-polar)

London dispersion,

dipole-dipole, H-

bonds

 Poor conductors  Low to moderate MP

CO2 (s), C12H22O11,

H2O (s)

Sucrose Carbon dioxide (dry ice)

Ice

Ionic Anions

and cations

Electrostatic attractions

 Hard & brittle  High MP  Poor conductors  Some solubility in H2O

NaCl, Ca(NO3)2

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Covalent (a.k.a.

covalent network)

Atoms bonded in

a covalent network

Covalent bonds

 Very hard  Very high MP  Generally insoluble  Variable conductivity

C (diamond & graphite)

SiO2 (quartz)

Ge, Si, SiC, BN

Diamond Graphite SiO2

Metallic

Metal cations in a

diffuse, delocalized

e- cloud

Metallic bonds

 Excellent conductors  Malleable  Ductile  High but wide range of MP

Cu, Al, Fe