Study Notes

The Kinetic Molecular Theory (KMT) explains the behavior of gases using 5 postulates:

1. Gases consist of tiny particles (atoms or molecules) that are in constant, RANDOM motion.

2. The particles themselves have NEGLIGIBLE VOLUME compared to the volume of the container. Gases are mostly empty space.

3. Gas particles exert NO ATTRACTIVE OR REPULSIVE FORCES on each other (no IMF between gas particles in ideal conditions).

4. Gas particles undergo PERFECTLY ELASTIC COLLISIONS — they bounce off each other and the container walls without losing energy. No kinetic energy is lost in these collisions.

5. The AVERAGE KINETIC ENERGY of gas particles is directly proportional to the TEMPERATURE (in Kelvin). Higher temperature = faster-moving particles = higher average KE.

Key concept: At the same temperature, all gas particles (regardless of mass) have the SAME average kinetic energy. Heavier particles move MORE SLOWLY; lighter particles move FASTER (same KE, different speeds).

Phase changes are physical changes involving transitions between solid, liquid, and gas states. They involve energy changes but NOT temperature changes during the transition.

Phase change names:

• Melting: solid → liquid (absorbs energy; endothermic)

• Freezing: liquid → solid (releases energy; exothermic)

• Boiling/Vaporization: liquid → gas (absorbs energy; endothermic)

• Condensation: gas → liquid (releases energy; exothermic)

• Sublimation: solid → gas directly (absorbs energy; endothermic)

• Deposition: gas → solid directly (releases energy; exothermic)

Endothermic phase changes: energy must be ABSORBED (melting, vaporization, sublimation)

Exothermic phase changes: energy is RELEASED (freezing, condensation, deposition)

Memory aid: Endothermic = energy INPUT needed to break bonds/forces = going from more ordered → less ordered state. Exothermic = energy RELEASED when IMF form = going from less ordered → more ordered.

A heating curve shows temperature vs. energy added for a substance as it goes from solid → liquid → gas.

Shape of the heating curve:

• Sloped line rising: solid being heated (temperature increases)

• Horizontal line (plateau): MELTING — solid and liquid coexist, temperature stays constant at the melting point

• Sloped line rising: liquid being heated (temperature increases)

• Horizontal line (plateau): BOILING — liquid and gas coexist, temperature stays constant at the boiling point

• Sloped line rising: gas being heated (temperature increases)

Reading a heating curve:

• The temperature where the first plateau occurs = melting point

• The temperature where the second plateau occurs = boiling point

• The SLOPE of each rising segment indicates specific heat capacity:

• Steeper slope = temperature rises faster = LOWER specific heat capacity

• Flatter slope = temperature rises slower = HIGHER specific heat capacity

Specific heat capacity (c): the amount of energy needed to raise 1 gram of a substance by 1°C. Water has an unusually HIGH specific heat (4.18 J/g·°C).

A phase diagram shows the physical state of a substance at different combinations of temperature and pressure.

Regions on a phase diagram:

• Solid region: low temperature, high pressure (usually top-left)

• Liquid region: moderate temperature and pressure (usually middle)

• Gas region: high temperature, low pressure (usually bottom-right)

• Supercritical fluid region: above the critical point

Key points:

• Triple point: the unique temperature AND pressure where all THREE phases coexist in equilibrium

• Critical point: the temperature and pressure beyond which gas and liquid phases cannot be distinguished — becomes a "supercritical fluid"

Reading a phase diagram:

• Find the temperature along the x-axis and pressure on the y-axis

• Read which region the point falls in → that's the physical state

• Freezing/melting point: where solid-liquid line crosses a given pressure

• Boiling point: where liquid-gas line crosses a given pressure

• Predict phase change: if you increase temperature at constant pressure, cross the boundary line = phase change occurs

Evaporation is the process where molecules at the SURFACE of a liquid escape into the gas phase at temperatures BELOW the boiling point.

For a molecule to evaporate, it must have enough kinetic energy to overcome the intermolecular forces holding it in the liquid.

Factors that increase evaporation rate:

• Higher temperature: more molecules have enough KE to escape

• Greater surface area: more molecules exposed at the surface

• Lower external pressure: less pressure suppressing the vapor phase

KE distribution graph (# particles vs. kinetic energy):

• Shows a curve: most particles have intermediate KE

• Higher temperature: curve shifts RIGHT (higher average KE)

• At higher temperature: MORE particles have enough energy to evaporate

• The shaded region above the minimum evaporation energy represents molecules that CAN evaporate