The Kinetic Molecular Theory

1. The Kinetic Molecular Theory

2. Review • Matter can exist in several states: • Solids have a definite shape and volume, and are not compressible. • Liquids have a definite volume but not a definite shape. They are not compressible. • Gases do not have a definite shape or volume. Gases are compressible. • Molecules are attracted to or repelled from each other by electrostatic forces called intermolecular forces. • Intermolecular forces include dipole-dipole interactions, hydrogen bonding, and London forces.

3. The Kinetic Molecular Theory • Kinetic molecular theory (KMT): • Explains the observed properties of matter in its different states. • Particles are in constant, rapid, random motion. • The speed of this motion depends on temperature. • Collisions between particles are elastic. • No kinetic energy lost.

4. Particles in Solids, Liquids, and Gases

5. The Kinetic Molecular Theory Right-click on the circle and select “Play” to see the animation. The motion of gas molecules is rapid, constant, and random.

6. The Kinetic Molecular Theory In a perfectly elastic collision, the kinetic energy of each molecule might change, but the total kinetic energy stays the same. Collisions between particles are elastic.

7. States of Matter • The state (solid, liquid, or gas) of a substance depends on three things: • Chemical identity • Temperature • Pressure • Chemical identity matters because: • Different substances have different degrees of intermolecular forces. • The molecules of O2 are small and non-polar. • Attracted to each other only by London forces. • O2 has a very low boiling point and is a gas at room temperature. • NaCl is a salt with very strong ionic bonds between ions. • NaCl has a very high melting point and is solid at room temperature.

8. States of Matter • In solids, attractive forces outweigh kinetic energy of particles. • Particles stick together in rigid position. • In liquids, attractive forces are balanced by kinetic energy. • Particles stick together but can move past each other. • In gases, kinetic energy outweighs attractive forces. • Particles fly rapidly past each other as if there are no attractive forces between them.

9. Pressure • Pressure = force / area

10. Pressure One gas molecule exerts a tiny force against the side of a balloon.

11. Pressure When you have a huge number of gas molecules colliding against the sides of a balloon, the force adds up. Force spread out over the inner surface of the balloon is pressure.

12. Units of Pressure • SI Unit: Pascals (Pa) • 1 Pa = 1 Newton / square meter • Other units: • atmospheres (atm) • millimeters of mercury (mmHg) • torr • pounds per square inch (psi)

13. Pressure Conversion Factors 1 atm = 760 mmHg 1 atm = 101.325 kPa

14. Standard Pressure • Standard Pressure = 1 atm • Also, 760 mmHg or 101.325 kPa • Standard pressure is the normal air pressure at sea level. • QUESTION: Why does air pressure decrease as you climb a mountain or ascend in an airplane?

15. Atmospheric Pressure As you move upward through the atmosphere, the density decreases. This is because most air molecules are held close to Earth’s surface by gravity. As the density decreases, there are fewer molecules colliding with surfaces; hence, less pressure.

16. Temperature • Temperature – measure of the average kinetic energy of a substance’s molecules. • The molecules in a hot object are moving faster than the molecules in a cold object on average!

17. Temperature

18. Temperature

19. Temperature

20. Temperature

21. Temperature • Absolute temp. is measured in Kelvins (K) • No degree symbol! • To convert: • from °C to K, add 273. • from K to °C, subtract 273. • EXAMPLE: • Water boils at 100°C, or 373 K. • Water freezes at 0°C, or 273 K. • Avg. human body temp. is 37°C, or 310 K.

22. Standard Temperature • Standard Temperature = 0°C. • Also, 273 K. • Absolute Zero = -273°C = 0 K. • Lowest temperature. • All molecular motion stops.

23. Phases and Phase Changes • Phase = state of matter Gas sublimation condensation deposition vaporization melting Solid Liquid freezing

24. Phase Diagrams • Phase Diagram - Shows the phases of a substance at various combinations of temperature and pressure.

25. Phase diagram of water

26. Phase diagram of CO2

27. States of Matter • Gas vs Vapor • gas – used to refer to a substance that is a gas under standard conditions. • Examples: O2, N2, H2, the noble gases, etc. • vapor – refers to the gaseous form of a substance that isn’t a gas under standard conditions. • Examples: water, alcohol, mercury, sodium

28. Evaporation • Evaporation – when liquid molecules escape as a gas below their boiling point. • In a closed container the gas molecules eventually return to the liquid. • Equilibrium is established. • In an open environment no equilibrium is established. • The liquid completely evaporates.

29. Boiling • The hotter a liquid is, the more vapor it produces. • Molecules move faster. • More molecules escape as vapor. • More vapor = more vapor pressure. • When a liquid’s vapor pressure = surrounding pressure, the liquid boils. • Boiling point – the temperature at which a liquid’s vapor pressure = surrounding pressure. • For water at 1 atm, b.p. = 100ºC (373 K) • When a liquid reaches its boiling point, bubbles of vapor form within the liquid itself.

30. Vapor Pressure vs. Temperature Vapor pressure of water