Chapter 12 & 13

Overview of Phases of Matter

• Three Phases of Matter: Gases, Liquids, and Solids.

Gases

• Definition: Substances that assume the shape and volume of their container.

• Example: Opening a bottle of perfume allows the vapor to fill the room.

• Compressibility: Gases are compressible; applying pressure can change their volume.

• Ideal Gas Law Connection: Discussed in previous courses.

Liquids

• Definition: Substances that assume the shape of the container but not its volume completely.

• Example: Pouring water into a glass—the water occupies only the volume filled.

• Compressibility: Liquids are not significantly compressible due to close particle arrangements.

Solids

• Definition: Substances retain their shape and volume regardless of the container.

• Example: A chair maintains its structure in a box during transport.

• Compressibility: Solids are also not compressible and maintain their shape under pressure.

Special Cases

• Sand: Although it appears to take the shape of a container, each grain retains its solid structure, classifying it as a solid.

• Temperature and Phase: Temperature greatly affects the state of matter, with lower temperatures typically favoring solid states.

Molecular Microscopy

• At Zero Kelvin: Molecular motion halts, but atomic sub-particle motion continues.

Intermolecular Forces

• Role: At low temperatures, intermolecular forces hold atoms/molecules tightly together, resulting in solid forms.

• As temperature increases, molecules gain energy, leading to phase changes (solid → liquid → gas).

Key Phase Changes

1. Melting: Solid to Liquid (Increasing temperature).

2. Vaporization: Liquid to Gas (Further increase in temperature).

3. Molar Mass Influence: Larger molar masses correlate to stronger intermolecular forces (London forces).

Molecular Geometry Review

Ammonia (NH3)

• Valence Electrons: 8 total (5 from N, 1 from each of 3 H).

• Geometry: Tetrahedral electron domain, trigonal pyramidal molecular geometry; results in polarity and permanent dipole moment.

Carbon Dioxide (CO2)

• Valence Electrons: 16 total (4 from C, 6 from each O).

• Geometry: Linear shape; symmetrical; nonpolar overall despite polar bonds.

Acetic Acid (CH3COOH)

• Structure: Recognized as a polar molecule with a functional carboxylic acid group; has a permanent dipole moment.

Intermolecular Forces

• Types:

  1. London Dispersion Forces: Weakest, occur in nonpolar molecules.

  2. Dipole-Dipole Forces: Occur in polar molecules; attraction between partial charges.

  3. Hydrogen Bonding: Stronger than dipole-dipole; occurs between electronegative atoms and hydrogen (N, O, F).

Vapor Pressure Influences

• Factors Affecting Vapor Pressure:

  • Increase in temperature leads to higher vapor pressure due to increased molecular escape into vapor.

  • Strong intermolecular forces lead to lower vapor pressure.

Phase Diagrams

Critical Points

• Critical Temperature and Pressure: Conditions beyond which a substance cannot be liquefied regardless of applied pressure.

• Triple Point: Condition where solid, liquid, and gas coexist.

Crystalline Solids

Cubic Unit Cells

• Types of Cubic Unit Cells:

  1. Primitive Cubic: 1 atom per unit cell defined by atoms at corners.

  2. Body-Centered Cubic: 2 atoms (adds a body atom in the center).

  3. Face-Centered Cubic: 4 atoms (adds atoms at the faces).

Relationships in Unit Cells

• Primitive Cubic Relationship: 2 radii = side length.

• Face-Centered Cubic Relationship: 4r = l√2 (hypotenuse of a right triangle formed by the face diagonal).

• Body-Centered Cubic Relationship: 4r = l√3 (formed through the body diagonal).

Heating Curves

• Understanding Heating Curves: Graphical representation that describes temperature response to added heat energy through phase changes.

  • Includes steps of cooling liquid water to ice over a range of temperatures.