1. Chemistry: The Central Science
Fourteenth Edition
Chapter 11
Liquids and Intermolecular Forces
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2. Intermolecular Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions (bonds) that hold compounds together.
Many physical properties reflect intermolecular forces, like boiling points, melting points, viscosity, surface tension, and capillary action.

3. States of Matter
The fundamental difference between states of matter is the strength of the intermolecular forces of attraction.
Stronger forces bring molecules closer together.
Kinetic energy keeps them apart and moving. Remember that average kinetic energy is related to temperature!

4. Properties and State of Matter
Table 11.1 Characteristic Properties of the States of Matter
Gas
Liquid
Solid
Assumes both volume and shape of its container
Assumes shape of portion of container it occupies
Retains own shape and volume
Expands to fill its container
Does not expand to fill its container
Is compressible
Is virtually incompressible
Flows readily
Does not flow
Diffusion within a gas occurs rapidly
Diffusion within a liquid occurs slowly
Diffusion within a solid occurs extremely slowly
*The atoms in a solid are able to vibrate in place. As the temperature of the solid increases, the vibrational motion increases.

5. Relative Strength of Attractions
Table 11.2 Melting and Boiling Points of Representative Substances
Force Holding Particles Together
Substance
Melting Point (K)
Boiling Point (K)
Chemical bonds
blank
Ionic bonds
Lithium fluoride (L i F)
1118
1949
Metallic bonds
Beryllium (B e)
1560
2742
Covalent bonds
Diamond (C)
3800
4300
Intermolecular forces
Dispersion forces
Nitrogen (N2) 63 77
Dipole–dipole interactions
Hydrogen chloride (H C l)
158
188
Hydrogen bonding
Hydrogen fluoride (H F)
190
293
Intermolecular attractions are weaker than bonds.
Hydrogen bonds are NOT chemical bonds.

6. Types of Intermolecular Force between Neutral Molecules
Weakest to strongest forces:
Dispersion forces (or London dispersion forces or induced dipole-induced dipole interactions)
Dipole–dipole forces
Hydrogen bonding (a special dipole–dipole force)
Note: The first two types are also referred to collectively as van der Waals forces.

7. Dispersion Forces
A nonpolar particle (helium atoms below) can be temporarily polarized to create dispersion force.
The tendency of an electron cloud to distort is called its polarizability.

8. Factors That Affect Amount of Dispersion Force in a Molecule
Number of electrons in an atom (more electrons, more dispersion force)
Size of atom or molecule/molecular weight
Shape of molecules with similar masses (more compact, less dispersion force)

9. Polarizability and Boiling Point
If something is easier to polarize, it has a lower boiling point.
Remember: This means less intermolecular force (smaller molecule or atom: lower molecular weight, smaller size, fewer electrons).

10. Dipole–Dipole Interactions (1 of 2)
Polar molecules have a more positive and a more negative end—a dipole
The oppositely charged ends attract each other.

11. Dipole–Dipole Interactions (2 of 2)
For molecules of approximately equal mass and size, the more polar the molecule, the higher its boiling point.

12. Which Have a Greater Effect: Dipole–Dipole Interactions or Dispersion Forces?
If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force.
If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

13. What Does This Graph Show Us?
In a group, the period 3/4/5 elements have higher boiling points as the group member gets larger.
What happens with the period 2 elements? For group 4A, the trend is continued. What about for the other groups?

14. Hydrogen Bonding
The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
We call these interactions hydrogen bonds.
A hydrogen bond is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.

15. What Forms Hydrogen Bonds?
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
These atoms interact with a nearly bare hydrogen nucleus (which contains one proton but NO inner electrons).

16. Ice Compared to Liquid Water
Hydrogen bonding makes the molecules farther apart in ice than in liquid water.

17. Ion–Dipole Interactions
Ion–dipole interactions are found in solutions of ions.
The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents, like water.

18. Determining Intermolecular Forces in a Substance
Type of Intermolecular interaction
Atoms Example: N e, A r
Nonpolar molecules Example:
B F3, C H4
Polar molecules
without O H, N H,
or H F groups Examples:
H C l, C H3C N
containing O H,
N H, or H F groups Examples:
H2O, N H3
Ionic solids
dissolved in
polar liquids Examples:
N a C l in H2O
Dispersion forces (0.1–30 kJ/mol) ✔
Dipole–dipole interactions
(2–15 kJ/mol)
blank
Hydrogen bonding (10–40 kJ/mol)
Ion–dipole interactions (>50 kJ/mol)
Note that ALL chemicals exhibit dispersion forces.
The strongest force dictates the extent of attractions between molecules.

19. Generalizations about Relative Strengths of Intermolecular Forces
When two molecules have comparable molar masses and shapes, dispersion forces are roughly equal.
When two molecules have very different molar masses and there is no H-bonding, dispersion force determines the substance with stronger attractions.

20. Liquid Properties Affected by Intermolecular Forces
Boiling point (previously discussed) and melting point
Viscosity
Surface tension
Capillary action

21. Viscosity (1 of 2) Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

22. Viscosity (2 of 2)
Table 11.3 Viscosities of a Series of Hydrocarbons at 20 degree Celsius

23. Surface Tension
Water acts as if it has a “skin” on it due to extra inward forces on its surface.
It causes water to “bead up” when in contact with nonpolar surfaces.
Those forces are called the surface tension.

24. Cohesion and Adhesion
Intermolecular forces that bind similar molecules to one another are called cohesive forces.
Intermolecular forces that bind a substance to a surface are called adhesive forces.
These forces are important in capillary action.

25. Capillary Action
The rise of liquids up narrow tubes is called capillary action.
Adhesive forces attract the liquid to the wall of the tube.
Cohesive forces attract the liquid to itself.
Water has stronger adhesive forces with glass; mercury has stronger cohesive forces with itself.

26. Phase Changes
Conversion from one state of matter to another is called a phase change.
Energy is either added or released in a phase change.
Phase changes: melting/freezing, vaporizing/condensing, subliming/depositing.

27. Energy Change and Change of State
The heat of fusion is the energy required to change a solid at its melting point to a liquid.
The heat of vaporization is the energy required to change a liquid at its boiling point to a gas.
The heat of sublimation is the energy required to change a solid directly to a gas.

28. Heating Curves
A graph of temperature vs. heat added is called a heating curve.
Within a phase, heat is the product of specific heat, sample mass, and temperature change.
The temperature of the substance does not rise during a phase change.
For the phase changes, the product of mass or moles and heat of fusion or vaporization is heat.

29. Supercritical Fluids (1 of 2)
Gases liquefy when pressure is applied.
The temperature beyond which a gas cannot be compressed is called its critical temperature. The pressure needed to compress the liquid at critical temperature is called critical pressure.
The state beyond this temperature is called a supercritical fluid.

30. Supercritical Fluids (2 of 2)
Table 11.5 Critical Temperatures and Pressures of Selected Substances
Substance
Critical Temperature (K)
Critical Pressure (a t m)
Nitrogen, N2
126.1
33.5
Argon, Ar
150.9
48.0
Oxygen, O2
154.4
49.7
Methane, C H4
190.0
45.4
Carbon dioxide, C O2
304.3
73.0
Phosphine, P H3
324.4
64.5
Propane, C H3C H2C H3
370.0
42.0
Hydrogen sulfide, H2S
373.5
88.9
Ammonia, N H3
405.6
111.5
Water, H2O
647.6
217.7

31. Vapor Pressure (1 of 4)
At any temperature, some liquid molecules have enough energy to escape the surface and become a gas.
As the temperature rises, the fraction of molecules that have enough energy to break free increases.

32. Vapor Pressure (2 of 4)
As more molecules escape the liquid, the pressure they exert increases.
The liquid and vapor reach a state of dynamic equilibrium: Liquid molecules evaporate and vapor molecules condense at the same rate.

33. Vapor Pressure (3 of 4)
The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.
The normal boiling point is the temperature at which its vapor pressure is 760 torr.

34. Vapor Pressure (4 of 4)
The natural log of the vapor pressure of a liquid is inversely proportional to its temperature.
This relationship is quantified in the Clausius Clapeyron equation.

35. Phase Diagram
A phase diagram is a graph showing states of matter under conditions of temperature and pressure.
It also shows changes of state and the triple point and critical point.

36. Phase Diagram of Water
Note the high critical temperature and critical pressure due to the strong van der Waals forces.
Unusual feature for water:
The slope of the melting curve is negative (as the pressure is increased, the melting point decreases).

37. Phase Diagram of Carbon Dioxide
Unusual features for carbon dioxide:
Cannot exist in the liquid state at pressures below atm (triple point).
C O2 sublimes at normal pressures.

38. Liquid Crystals (1 of 2)
Some substances do not go directly from the solid state to the liquid state.
In this intermediate state, liquid crystals have some traits of solids and some of liquids.
Molecules in liquid crystals have some degree of order.

39. Liquid Crystals (2 of 2)
In nematic liquid crystals, molecules are ordered in only one dimension, along the long axis.
In smectic liquid crystals, molecules are ordered in two dimensions, along the long axis and in layers.
In cholesteric liquid crystals, nematic-like crystals are layered at angles to each other.

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