AP/IB Chemistry Notes Name:
Chapter 10 - Liquids and Solids
I. Intermolecular Forces
A. Intramolecular vs. Intermolecular Forces
1. intramolecular forces
a.
b. molecule
c. or
2. intermolecular forces
a. molecules
b. causes solids or liquids (condensed states of matter) to form as molecules bond together
c. determines many important properties of substances
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II. Types of Intermolecular Forces:
A. Ion-dipole force
1. the attraction of an ion and a dipole
B. Dipole - dipole attractions ( molecules)
1. Molecules with dipoles line up so that the of one molecule is close to the of another molecule.
2. The attraction as the distance between molecules
3. Much than or forces (the dipole-dipole force is typically about of the force of an ionic or covalent bond)
4. But does explain why polar liquids are more soluble in than in non-polar liquids
a. It takes mL of H2O to dissolve 1 mL of CCl4
b. It takes mL of H2O to dissolve 1 mL of CH2Cl2
C. London - dispersion forces ( molecules)
1. Exist in non-polar molecules and noble gas atoms
2. Instantaneous can be produced in non-polar molecules when electrons are .
3. When an instantaneous occurs in one molecule, dipoles are in the neighboring molecules. (figure 10.5)
4. Weaker than attractions and .
5. Increase with .
a. Larger molecules have London-dispersion forces. (more “polarizable”)
b. Larger molecules have , creating a greater opportunity for .
D. Hydrogen Bonding
1. Hydrogen bond is an especially strong , as shown by the trend in boiling points of polar molecules.
2. H-bonding is observed for:
3. Conditions for occurrence:
a. H attached to a small, element in one molecule
b. Small, highly electronegative element with in the other molecule
4. Observed for the elements:
5. Which of the following molecules will hydrogen-bond in the pure substance?
a. H2O
b. H2Se
c. HF
d. HBr
e. NH3
f. PF3
H-Bonding for liquid water
6. H points at the in the other molecule
7. In liquid water, each water molecule is surrounded by an average of
other water molecules; structure is not rigid.
8. than covalent bond.
9. Average of 4 hydrogen bonds in liquid water
10. Fluoride ion is hydrogen-bonded to water in solution
11. Molecules hydrogen-bond to themselves or to other molecules.
Structure of Ice
12. The water molecules in ice are fixed into a arrangement as a result of hydrogen bonding. Open structure makes ice than water.
13. The open structure of ice leaves channels of through the crystals.
AP/IB Chemistry Notes Name:
10.1 Intermolecular Forces Summary
Identifying Types of molecular forces
What types of intermolecular forces are observed for each of the following molecules?
(A molecule may have more than one.)
H2O HF
HBr NH3
PF3 CH3OH
F2 CO
CO2 N2
Strength of Intermolecular forces
14. Intermolecular forces generally increase in strength as:
a.
15. The forces are . All molecules have London forces. Polar molecules have both London and dipole-dipole forces. ...
16. Which member of each pair has the larger intermolecular forces? (we will learn later that this affects things such as boiling point, heat of vaporization)
CH3OH, CH3SH CO, HF
F2, Kr CO2, NH3
F2, CO N2, NH3
AP/IB Chemistry Notes Name:
Chapter 10 - Liquids and Solids
III. Liquid State Phenomena
A. Structural models of liquids
1. More complex than models for solids or gases for two reasons
a. Liquids have strong forces
b. Liquids have significant
A. Surface tension - the of a liquid to increase its surface area.
1. For the surface area of a liquid to increase, molecules would have to move up to the surface.
a. This would require internal molecules to from their surrounding molecules, going the intermolecular forces.
2. An distribution of forces exists on molecules
a. Molecules on the surface only experience intermolecular attractions with molecules of them.
b. Molecules below the surface experience intermolecular attractions with molecules .
c. This causes molecules on the surface to be , giving the surface a .
3. Liquids with intermolecular forces have a surface tension.
4. Would you expect a polar liquid to have a greater surface tension than a nonpolar liquid? Why or why not?
B. Capillary action - the of a liquid in a .
1. This occurs when the molecules in the container have .
2. The polar bonds in the container attract the liquid, causing the liquid to try to creep up the sides of the container, which stretches the surface of the liquid.
3. As the liquid tries to balance the attraction (cohesive forces) and the attraction between the (adhesive forces), the liquid the tube.
4. We can use this property to explain the shape of the meniscus formed by a liquid in a tube.
a. The meniscus of water is because the attractions between the molecules and the molecules are greater than the attractions molecules.
5. What would you expect the meniscus to look like for a liquid in which the internal (liquid to liquid) attractions are stronger than the attractions between the liquid and the container?
C. Viscosity - a liquid's
1. Viscosity can be compared to the " " of a liquid.
a. Thicker liquids are .
b. Example: Maple syrup has a greater viscosity than water.
2. Intermolecular forces and molecular complexity both contribute to the viscosity of a liquid.
a. As intermolecular forces , viscosity .
b. As molecular complexity , viscosity .
3. Why do you think that viscosity increases when the intermolecular forces or molecular complexity increases?
D. In what way do you think that viscosity, capillary action, and surface tension are related? In other words, if a liquid had a high viscosity, what would you predict about its surface tension and capillary action?
AP/IB Chemistry Notes Name:
10.8 Vapor Pressure & Phase Diagrams
IV. Phase Changes
A. Physical states of a substance can co-exist under a variety of conditions of pressure and temperature.
B. Phases: different forms (gas, liquid, solid, etc.) of a substance that co-exist in a heterogeneous system.
C. Transitions between phases are called phase changes
a. evaporation:
b. melting:
c. sublimation:
D. What are the phase changes?
Label on the diagram:
E. Heat of Vaporization
a. ΔHvap = energy needed to of liquid at constant temperature
b. Energy used to during evaporation
c. Larger molecules have because of higher London forces
d. > if molecular size is similar
e. H-bonded >
F. Heat of Fusion
a. Generally heat of fusion (enthalpy of fusion) is heat of vaporization:
b. it takes more energy to molecules, than them.
G. Freezing
a. Cooling liquids their kinetic energy
b. When intermolecular forces become than kinetic energy, the liquid and becomes solid.
c. Freezing point = temperature at which solid and liquid are in a state of .
d. Normal freezing point = f.p. at pressure of .
V. Vapor Pressure
A. Evaporation: loss of , so the liquid cools (unless energy is supplied) (The process is endothermic.)
B. Evaporative cooling, perspiration, alcohol bath, canvas water bags, wind chill factor
C. Evaporation:
a. Open container: evaporates completely
b. Closed container: reach a state of equilibrium
D. Equilibrium:
E. P at equilibrium =
F. When Pvap = Patm, T =
G. When Pvap = 1 atm, T =
VI. Phase Diagrams
B. Phase diagram: plot of summarizing all between phases.
C. Given a temperature and pressure, phase diagrams tell us which phase will exist.
a. lines:
b. areas:
c. triple point (confluence of 3 lines):
D. Features of a phase diagram:
a. Triple point: temperature and pressure at which all three phases are in .
b. Vapor-pressure curve: generally as pressure increases, temperature .
c. Critical point: critical temperature and pressure for the gas.
d. Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid.
e. Normal melting point: melting point at 1 atm.
IIII. The Wide World of Solids
A. A concept map to organize how solids are classified.
B. Amorphous solids
1. Highly
2. Described as a
3. Example: Glass
C. Crystalline solids
1. Highly arrangement
2. Represented as a
3. Unit cell - of a lattice
a. Three types of unit cells (figure 10.9)
1. - a cube with atoms or molecules at each of the corners
2. - a cube with atoms or molecules at each of the corners and one atom or molecule in the center of the cube.
3. - a cube with atoms or molecules at each of the corners and atoms or molecules in the center of each side of the cube
4. Structures determined by using X-ray diffraction
a. This technology was helpful in determining the structure of DNA!
5. Three main types of crystalline solids exist: ionic, molecular, and atomic.
D. Ionic solids
1. Characteristics
a. Stable
b. High
c. Brittle
d. Held together by
e. when solid
f. What do you think would happen to the conductivity of an ionic solid if it were melted or dissolved in water? Why?
2. Structural models
a. Similar to that of
b. Ions represented as
c. Ions are packed and arranged in a way to maximize attraction and
d. Example: NaCl (figure 10.35)
1. Chlorine ions are arranged in a face-centered cubic closest packed structure
2. Sodium ions fill in the spaces in between the Chlorine ions
E. Molecular solids
1. Lattice positions occupied by
2. Characteristics
a. Strong covalent bonding the molecules.
b. Molecules held to each other by (weaker).
c. Generally have
d. Generally of heat and electric current
e. Others vary with the type of molecule forming the solid
3. Molecules are held together by
a. The type of forces also varies with the type of molecule
1. Polar molecules are held together by
2. Nonpolar molecules are held together by .
b. Intermolecular forces are when the molecules are polar
F. Atomic solids
1. Consist of at each of the lattice points in the crystal
2. Can be divided into Metallic, Network, and Group 8A atomic solids
G. Metallic solids
1. Characteristics of metallic solids
a. High conductivity
b. High conductivity
c. Malleable
d. Ductile
e. Range from low melting point and soft to high melting point and brittle
f. Strong, covalent bonding
2. Structure of Metallic Crystals: The Closest Packing Model
a. Metal atoms are represented as
b. In a metallic solid, atoms are arranged in , packed as close together as possible, in a way that the space between atoms is minimized.
c. There are two common ways that metal atoms can be packed together
1. Hexagonal closest packed structure (figure 10.14)
a) Atoms arranged in the form
b) Creates a unit cell
c) Examples: Magnesium and Zinc
2. Cubic closest packed structure (figure 10.15)
a) Atoms arranged in the form
b) Creates a unit cell
c) Examples: Silver and Copper
d. Not all metallic solids have one of these two structures
a) e.g. alkali metals
b) unit cell
c) eight nearest neighbors (less dense)
3. Bonding Models for Metallic Solids
a. Electron Sea Model
1. Represented as surrounded by a sea of
2. Similar to Thomson's Plum Pudding model of the atom.
b. Molecular Orbital Model
1. The valence orbitals of the metals atoms hybridize in a way that allows valence electrons to move throughout the entire crystal
4. Metal Alloys
a. Alloy - a substance containing a and having metallic properties
b. Two general types of alloys exist
1. alloys - within a metal crystal, metal atoms are removed and replaced with atoms of similar size
a) Example: Brass (figure 10.21a)
2. alloys - small atoms fill in the holes in the metal crystal
b) Example: Steel (figure 10.21b)
c) Can make the metal stronger by creating directional bonds between the metal atoms and the smaller atoms
H. Network Atomic Solids
1. Characteristics
a. Brittle
b. Poor (thermal and electrical)
c. Very
d. Strong, covalent bonds
2. Often referred to as " "
3. Have type structures (figure 10.22)
a. Carbon in diamond form is arranged in tetrahadrals connected together
4. Examples: Carbon and Silicon (see Section 10.5)
IV. Changes of State
A. Changing into a Gas: Fun with Vapor Pressure
1. (evaporation) - molecules of a liquid escape the liquid's surface and form a gas
2. ( DHvap) - the energy that is required to vaporize one mole of liquid at one atmosphere of pressure
3. - vapor (gas) molecules going back into the liquid phase
4. - solid molecules escaping into the gas phase without going through the liquid phase
a. Can you think of an example of sublimation?
5. - the pressure that is exerted by a vapor (gas), in a closed system, measured when the equals the
a. When the vapor pressure is , a large number of the liquid molecules are taking part in the equilibrium
b. As the vapor pressure , the rate of evaporation .
1. More of the molecules are in the gas phase
c. The size of intermolecular forces has the greatest effect on vapor pressure
1. As the intermolecular forces , the vapor pressure
2. Why do you think that increasing intermolecular forces decreases the vapor pressure?
d. Molar mass also affects the vapor pressure
1. As the molar mass , the vapor pressure
e. Vapor pressure increases with temperature
1. As the temperature , more molecules will have enough energy to escape.
2. The relationship between vapor pressure and temperature is non-linear and can be represented by the following equation
where Pvap = vapor pressure, Hvap = heat of vaporization, T = temperature in K R = universal gas constant, C = constant
3. This equation can be rearranged to compare vapor pressure at two different temperatures
B. The Heating Curve
1. Diagram (figure 10.42)
2. During a phase change, the temperature remains even though heat is being added continually
a. If heat is still being added, why does the temperature not increase during a phase change?