1. Chapter #12 States of Matter Inter-particle Forces

2. Chapter #12 Intermolecular Forces (IMF) Topics Molecular interactions
Properties of Liquids and Solids
IMF Force Properties
Evaporation and Condensation
Melting Freezing and Sublimation
Types of IMF
Types of Crystalline Solids

3. 12.1 Intramolecular Forces
“Within” the molecule and are called covalent and ionic bonds.
Molecules are formed by sharing electrons between the atoms.
Covalent bonds hold the atoms of a molecule together.
Ionic bonds hold oppositely charged ions together in a formula unit

4. 12.1 Intermolecular Forces
Forces that occur between molecules.
Intramolecular (ionic and covalent) bonds are stronger than intermolecular forces.

5. 12.6 Intermolecular Forces
There are two kinds of intermolecular forces that occur between molecules.
Dipole–dipole forces (between polar molecules)
Hydrogen bonding (between hydrogen of one polar molecule and O, N, F of another polar molecule)
London dispersion forces (between nonpolar molecules)

6. 12.6 Dipole–Dipole Attraction
Polar molecules contain different
atoms which have different
electronegativity's and can be
thought to be like a small bar
magnet. Oppositely charged ends attract and like ends repel. Molecules organize themselves to maximize the attractive forces
and minimize repulsive forces
giving molecules unique shapes.

7. 12.6 Dipole-Dipole Forces
Dipole moment – molecules with polar bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge.
Molecules with dipole moments can attract each other electrostatically. They line up so that the positive and negative ends are close to each other.
Only about 1% as strong as covalent or ionic bonds.

8. 12.6 Hydrogen Bonding Strong dipole-dipole forces.
Hydrogen is bound to a highly electronegative atom – nitrogen, oxygen, or fluorine.

9. 12.6 Hydrogen Bonding in Water
Blue dotted lines are the intermolecular forces between the water molecules.

10. 12.6 Hydrogen Bonding
Affects physical properties
Boiling point

11. 12.6 London Dispersion Forces
Instantaneous dipole that occurs accidentally in a given atom induces a similar dipole in a neighboring atom.
Significant in large atoms/molecules.
Occurs in all molecules, including nonpolar ones.

12. 12.6 London Dispersion Forces
Become stronger as the sizes of atoms or molecules increase.

13. 12.6 London Dispersion Forces
Nonpolar Molecules

14. Ion Dipole Forces
Ion Dipole forces are typically found in aqueous solutions of ionic compounds, such as salt water. In the figure below the nonbonding electrons found on the oxygen atom are strongly attracted to the sodium ion.

15. Overview of Particle Forces
Type of Force Relative Strength Present in Examples
Dispersion Force
(London Force)
Weak, increasing with size
atoms/molecules
H2 (g)
Polar molecules
HCl
Dipole-dipole force
moderate
Molecules with hydrogen bonded to N,O, or F
Hydrogen bond
Strong
HF, NH3, H2O
Ion-dipole
Very strong
Mixtures of ionic and polar compounds
Na+,Cl-, and H2O

16. 12.2 The Liquid State
In the liquid state the particles are randomly arranged (like in a gas).
The particles are closer to each other than in gases so the density of liquids is greater than that of gases, thus attractive forces between liquid particles is stronger than in gases.
Liquids adopt the shape of the container into which they are placed.
Unlike gases liquids have a fixed volume and density. This is because they have greater attractive forces between particles holding them close together.

17. 12.2 The Liquid State
As the particles in liquids are very close to one another they have small compressibility.
When particles of a liquid are heated the particles will move around more rapidly.
As they have many close neighbors they may only travel a short distant before undergoing a collision and bouncing back in the opposite direction. For this reason liquids have little thermal expansion.

18. 12.2 Gas/Liquid Comparison
Liquid state
Gaseous state
Particles are close together
Not held in fixed positions
Take the shape of container
Have fixed volume
Little compressability
Small thermal expansion
Particles are far apart
Completely fill container
Easily compressed
Moderate thermal expansion

19. 12.2 Solid State
In the solid state particles are held in fixed lattice positions.
Cohesive forces are much more dominant for solids than dispersive forces.
Unlike gases solids have fixed shape, volume and density.
The particles may only move a small amount around their fixed positions so solids have little thermal expansion.
In solids the particles are close together and so they have high density.

20. 12.2 Solid State Strong cohesive forces
Particles in fixed lattice positions
Constant shape
Constant density
Constant volume
Minimal compressibility
Little thermal expansion

21. 12.3 Surface Tension

22. Viscosity
Viscosity is the resistance to flow. Viscose liquids slowly flow such as syrup, while non viscose liquids flow rapidly, such as water. The stronger the antiparticle forces the more viscose the liquid

23. 12.3 Cohesive and Adhesive Forces

24. Beading and Wetting
Beading is desired for cars, while wetting is desired for the dishwasher. Beading in a dishwasher produces spots on the glassware. Beading (strong cohesive forces) and wetting (strong adhesive forces) are illustrated below.

25. 12.4 Changes in States of Matter
When matter takes energy from its surroundings (endothermic processes) the kinetic energy of the particles increases resulting in greater dispersive forces and the particles moving away from each other.
i.e. Processes in which particles move away from each other (solid to liquid change of state) are endothermic.

26. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other. i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
H2O (s) H2O (l)
H2O (g)

27. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
H2O (s) H2O (l)
H2O (g)
melting

28. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
H2O (s) H2O (l)
H2O (g)
melting

29. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
H2O (s) H2O (l)
H2O (g)
melting
evaporation

30. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
H2O (s) H2O (l)
H2O (g)
melting
evaporation
sublimation

31. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
condensation
H2O (s) H2O (l)
H2O (g)
melting
evaporation
sublimation

32. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
condensation
H2O (s) H2O (l)
H2O (g)
melting
evaporation
sublimation

33. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
freezing
condensation
H2O (s) H2O (l)
H2O (g)
melting
evaporation
sublimation

34. 12.4 Changes in States of Matter
When matter releases energy to its surroundings (exothermic processes) the kinetic energy of the particles decreases resulting in greater cohesive forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each other (liquid to solid change of state) are exothermic.
deposition
freezing
condensation
H2O (s) H2O (l)
H2O (g)
melting
evaporation
sublimation

35. 12.4 Changes in States of Matter
The specific heat of a solid, liquid, or gas is the amount of energy required to increase the temperature of one gram by one degree C.
The heat of fusion of a solid is the amount of heat required to change one gram of solid to liquid.
The heat of vaporization of a liquid is the amount of heat required to change one gram of liquid to a gas.

37. Compound Specific Heat

38. Heating Curve

39. 12.4 Sample Heating Problem
Calculate the energy in Kj required to change 5.00g of ice into water at 66°C.
First find energy to melt ice using heat of fusion.
5.00 g
333 j Kj
= Kj g
103 j
Now find the energy to heat water from 0°C to 66°C
4.184 j
5.00 g
66 °C Kj
= Kj
g-°C
103 j
1.665 Kj
Kj = 3.05 Kj

40. 12.4 Evaporation
Evaporation is the name of the process by which a liquid becomes a gas.
Considering our previous discussion would you expect evaporation to be endothermic or exothermic?
Evaporation takes place from the surface of a liquid.
How do you expect the rate of evaporation to be affected by the surface area of the liquid?

41. 12.4 Vapor Pressure We can define the vapor pressure of a liquid as:
“the pressure exerted by a vapor that is in equilibrium with its liquid.”
This is a little confusing so lets take sometime to explain.

42. 12.4 Vapor Pressure
If we place a liquid in a sealed conatiner with some empty space above the liquid initially there will be no vapor or gas above that liquid.
Those molecules at the surface of the liquid with sufficient energy will leave the liquid and enter the gas phase. Some of the vapor molecules will strike the surface of the liquid and return to the liquid phase.

43. 12.4 Vapor Pressure
When the rate at which the liquid is entering the gas phase equals the rate at which the vapor is returning to the liquid phase we say the system is at equilibrium. After this time the liquid level will remain constant. The pressure exerted by the vapor at this time is called the vapor pressure.

44. 12.4 Vapor Pressure
The vapor pressure of a liquid decreases with molecular mass.
The vapor pressure of a increases with temperature.
The vapor pressure of a liquid depends upon the chemical nature of the liquid.
Those molecules that have strong intermolecular attractive forces have lower vapor pressures than expected for their molecular mass.

45. 12.4 Vapor Pressure Lets consider some examples:
At 20oC H2O (MW = 18 gmol-1) has a vapor pressure of 17.5 torr !!
This due to strong hydrogen bonds between water molecules.

46. 12.4 Vapor Pressure
As we increase the temperature the vapor pressure of a liquid increases.
The temperature at which the vapor pressure equals the external pressure (atmospheric pressure) is called the boiling point.
Bubbles of vapor with atmospheric pressure may form anywhere in the liquid and rise to the surface at the boiling point of a liquid.
What is the effect of lowering the atmospheric pressure on the boiling point?

47. 12.5 Melting and Boiling Points
In general, the stronger the intermolecular forces, the higher the melting and boiling points.

48. 12.4 Vapor Pressure
As we increase the temperature the vapor pressure of a liquid increases.
The temperature at which the vapor pressure equals the external pressure (atmospheric pressure) is called the boiling point.
Bubbles of vapor with atmospheric pressure may form anywhere in the liquid and rise to the surface at the boiling point of a liquid.
What is the effect of lowering the atmospheric pressure on the boiling point? Lowering of the boiling point, thus food takes longer to cook at higher elevations.

49. 12.7 Types of Crystalline Solids
Molecular Solids: Solids made of individual molecules attached by interpartical forces.
Examples: ice, dry ice, and sugar
Properties: Low melting, Non conductors of heat and electricity, and brittle.
Ice

50. 12.7 Types of Crystalline Solids
Covalent Network Solids: Atoms or molecules held together by covalent bonds.
Examples: Diamond, Bucky balls, Quartz
Properties: High melting, nonconductors of heat and electricity and brittle
Quartz
Bucky Balls

51. 12.7 Types of Crystalline Solids
Ionic Solids: Solids held together by ionic bonds. They are never ending arrays of oppositely charged ions.
Examples: Salt (sodium chloride), Copper (II) nitrate
Properties: High melting, brittle, nonconductors of electricity.
Sodium Chloride

52. 12.7 Types of Crystalline Solids
Atomic Solids: Solids consisting of individual atoms.
Examples: Selenium, sulfur, and Neon
Properties: Low melting and boiling points, non conductors of electricity, and brittle.
sulfur

53. 12.7 Types of Crystalline Solids
Metallic Solids: Solids composed of metals
Examples: Copper, Iron, Silver and Gold
Properties: Variable melting and boiling points. Eg. Mercury melts at – 39°C, while Tungsten melts at 3422°C
Electron Sea Model

54. End Chapter #12