2009 Chemistry I: Modern Chemistry Holt, Rinehart, & Winston
Chapter 10 – States of Matter, pp 329-351
Vocabulary
1 amorphous solid; 2 boiling; 3 boiling point; 4 capillary action; 5 condensation; 6 critical point;
7 critical pressure; 8 critical temperature; 9 crystal; 10 crystalline solids; 11 crystal structure;
12 deposition; 13 diffusion; 14 effusion; 15 elastic collision; 16 equilibrium; 17 equilibrium vapor pressure; 18 evaporation; 19 fluid; 20 freezing point; 21 ideal gas; 22 kinetic-molecular theory;
23 melting; 24 melting point; 25 molar enthalpy of fusion; 26 molar enthalpy of vaporization;
27 phase; 28 phase diagram; 29 real gas; 30 sublimation; 31 supercooled liquids; 32 surface tension
33 triple point; 34 unit cell; 35 vaporization; 36 volatile liquids
1 solids in which the particles are arranged randomly
2 a change of a liquid to bubbles of vapor that appear throughout the liquid as well as at its surface
3 temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure
4 attraction of surface of a liquid to the surface of a solid
5 process by which a gas changes to a liquid
6 the critical temperature & critical pressure of a substance
7 the lowest pressure at which the substance can exist as a liquid at the critical temperature
8 temperature above which a substance cannot exist in the liquid state
9 substance in which particles are arranged in an orderly, geometric, repeating pattern
10 solids consisting of crystals
11 the total 3-dimensional arrangement of particles of a crystal
12 change of state from a gas directly to a solid
13 spontaneous mixing of the particles of 2 substances caused by their random motion
14 process by which gas particles pass through a tiny opening
15 a collision in which there is no net loss of total kinetic energy
16 dynamic condition in which 2 opposing changes occur at equal rates in a closed system
17 pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature
18 process by which particles escape from the surface of a non-boiling liquid & enter the gas state
19 substance that can flow and will take the shape of its container
20 temperature/ pressure at which the liquid & solid of a substance are at equilibrium
21 a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory
22 an explanation for the properties & behavior of solids, liquids, & gases in terms of energy & forces
23 physical change of a solid to a liquid by addition of thermal energy
24 the temperature/ pressure at which a solid becomes a liquid
25 amount of thermal energy required to melt 1 mole of solid at the solid’s melting point
1 amorphous solid; 2 boiling; 3 boiling point; 4 capillary action; 5 condensation; 6 critical point;
7 critical pressure; 8 critical temperature; 9 crystal; 10 crystalline solids; 11 crystal structure;
12 deposition; 13 diffusion; 14 effusion; 15 elastic collision; 16 equilibrium; 17 equilibrium vapor pressure; 18 evaporation; 19 fluid; 20 freezing point; 21 ideal gas; 22 kinetic-molecular theory;
23 melting; 24 melting point; 25 molar enthalpy of fusion; 26 molar enthalpy of vaporization;
27 phase; 28 phase diagram; 29 real gas; 30 sublimation; 31 supercooled liquids; 32 surface tension
33 triple point; 34 unit cell; 35 vaporization; 36 volatile liquids
26 amount of thermal energy needed to vaporize 1 mol of liquid at its boiling point at constant pressure
27 any part of a system that has uniform composition and properties
28 graph of pressure vs. temperature showing the conditions under which phases of a substance exist
29 gas that does not behave completely according to the assumptions of the kinetic-molecular theory
30 change of state from a solid directly to a gas
31 substances that retain liquid properties even at temperatures at which they appear to be solid
32 force tending to pull adjacent parts of a liquid’s surface together, decreasing surface area maximally
33 temperature & pressure at which the solid, liquid, & vapor of a substance can coexist at equilibrium
34 the smallest portion of a crystal lattice that shows the 3-D pattern of the entire lattice
35 process by which a liquid or solid changes to a gas
36 liquids that evaporate readily
Main Ideas
I. Kinetic-molecular Theory of Matter pp 329-332
A. Kinetic-molecular theory of gases
1. Assumptions of kinetic-molecular theory
a. Gases consist of large numbers of tiny particles that are far apart relative to their size.
b. Collisions between gas particles & particles of container walls are perfectly elastic.
1. Total kinetic energy of particles remains constant at a constant temperature.
c. Gas particles are in constant, rapid, random motion, therefore possessing kinetic energy.
d. There are no forces of attraction between gas particles.
e. The temperature of a gas depends on the average kinetic energy of the gas particles.
1. KE = ½ mv2 where m = mass & v = velocity
2. All gases at same temperature have same kinetic energy, so lighter gases move faster.
2. Applies only to ideal gases
3. Real gases approach “ideal” at low pressures & high temperatures
B. Kinetic-molecular theory & the Nature of Gases
1. Expansion
a. Gases expand to fill container because particles move in all directions without significant
attraction between particles.
2. Fluidity
a. Due to insignificant attraction between particles, particles slide past one another like liquids
b. Because both gases & liquids flow, they’re both referred to as fluids.
3. Low Density
a. Because gas particles are so far apart from each other, density is 1/1000th that of liquid state
4. Compressibility
a. Because gas particles are so far apart from each other, volume can be greatly decreased.
5. Diffusion & Effusion
a. Because gas particles are in constant random motion, they will mix with other particles
b. Rates of effusion are directly proportional to velocities of their particles
1. Smaller masses effuse more quickly.
c. Graham’s law of effusion
1. Rate is inversely proportional to the square root of gas’s molar mass.
2. RateA / RateB = √molar massB / √molar massA
6. Deviation of Real Gases from Ideal Behavior
a. Particles occupy real space.
b. Particles exert attractive forces on each other.
c. Noble gases, especially He & Ne, are most likely to behave “ideally”.
d. Diatomic gases, such as H2 & N2, are nonpolar & also “ideal” under certain conditions.
e. The more polar a gas/ vapor is, such as H2O & NH3, the less “ideal” it is.
II. Liquids pp 333-336
A. Properties of Liquids & the Kinetic-molecular Theory
1. Attraction between particles is more effective than those between gas particles.
2. Attraction is caused by intermolecular forces.
a. Dipole-dipole forces
b. London dispersion forces
c. Hydrogen bonding
3. Particles are not bound together in fixed positions & thus a liquid is considered a fluid.
a. Most liquids flow downhill due to gravity, BUT liquid He near 0 Kelvin flows uphill!!!
B. Properties
1. Relatively High Density
a. At normal atmospheric pressure, most liquids are 100x denser than their gas state.
b. Most liquids are only ~ 10% less dense than their solid state.
· Water is an exception. Water’s greatest density is at 4o C.
c. Liquids differ greatly in density, thus having the ability to separate into layers.
2. Relative Incompressibility
a. Because liquid particles are closer together than are gases’, they are less compressible.
b. Liquid water at 1000 atm experiences only a 4% decrease in volume.
c. Liquid transmits pressure equally in all directions. (Pascal’s principle)
3. Ability to Diffuse
a. Any liquid gradually diffuses throughout any other liquid in which it can dissolve.
b. Liquids diffuse more slowly than gases because forces are greater & spaces are less
intermolecularly.
c. Increase in temperature (increase in average kinetic energy) increases rate of diffusion
4. Surface Tension
a. Result of intermolecular attractive forces.
b. Water has higher surface tension than most liquids due to H-bonding.
c. Liquid droplets take on spherical shape.
· A sphere has the smallest possible surface area for a given volume.
d. Capillary action pulls surface of liquid upward in a narrow tube against force of gravity.
· Liquid rises until attractive forces between liquid & surface = weight of liquid.
· Partially responsible for transport of water from roots to leaves in xylem of plants.
· Responsible for concave surface of meniscus in test tubes & graduated cylinders.
5. Evaporation & Boiling
a. Evaporation is a form of vaporization.
b. Evaporation occurs because particles of a liquid have different kinetic energies.
c. If kinetic energy exceeds intermolecular forces, particles escape into vapor phase.
d. Evaporation is critical to biological systems.
· Removes fresh water from ocean surface to concentrate salt & provide rain/ snow.
· Perspiration keeps organisms cool as the water absorbs the heat.
6. Formation of Solids
a. As temperatures decrease, particle motion decreases & intermolecular attractive forces pull
particles into a more orderly arrangement.
b. Solidification, i.e., freezing, occurs at different temperatures for different substances.
· H2O freezes at 0oC
· Ethanol freezes at -114oC.
· Paraffin is solid at room temperature.
III. Solids pp 337-341
A. Properties of Solids & the Kinetic-molecular Theory
1. Interparticle attractions, dipole-dipole, London dispersion, H bonding, are greatest in solids.
2. Have only vibrational motion around fixed points.
3. Have 2 types of solids: crystalline or amorphous
B. Properties
1. Definite Shape & Volume
a. Maintain without a container.
b. Crystalline solids are geometrically regular vs. amorphous solids without geometric shapes.
c. Volume changes only slightly with a change in temperature or pressure
2. Definite Melting Point
a. Attractive forces holding particles together are overcome by increase in their kinetic energy
· Crystalline solids have definite melting points.
· Solids covalently bound together have lower melting points than those of ionic bonds
· Amorphous solids (glass & plastic) have no definite melting point due to random arrangement of their particles.
3. High Density & Incompressibility
a. Particles of solids are more closely packed together than those of liquids or gases, usually.
· Solid H2 has lowest density, 1/320 that of the densest element, Os
· Cork may seem compressible, but it’s really spaces between the solid that are reduced.
4. Low Rate of Diffusion
a. Metals, such as Zn next to Cu, can diffuse into one another, but it’s a million times slower.
C. Crystalline Solids
1. Can exist as either single crystals or groups of crystals fused together.
a. The arrangement of particles can be represented by a coordinate system called a lattice.
b. A crystal & its unit cell can have 1 of 7 types of symmetry.
· cubic (fluorite)
· tetragonal (chalcopyrite)
· hexagonal (emerald)
· trigonal (calcite)
· orthorhombic (aragonite)
· monoclinic (azurite)
· triclinic (rhodonite)