1. States of Matter A Matter of Kinetic Energy

2. Types of States of Matter Solid Liquid Gas Plasma Beam BEC, or Bose-Einstein Condensate – Zero State of Matter – Most Dense

3. Changes of State

4. Kinetic Energy (kelvins & paschals)

5. chemwiki.ucdavis.edu/Physical_Chemistry/Physical_Properties_ of_Matter/Supercritical_Fluids Supercritical fluids are useful in science today – extraction of floral fragrance – the process of creating decaffeinated coffee – food science and functional food ingredients – pharmaceuticals, cosmetics, polymers, powders, bio- and functional materials – nano-systems, natural products, biotechnology, fossil & biofuels, microelectronics & environment (Bottini 133).

6. www.engineeringtoolbox.com/vapor-steam-d_609 Superheated Vapor When the temperature is higher than the boiling point @ a given pressure. Vapor cannot exist in contact with the fluid, nor contain fluid particles. Increase in pressure or decrease in temperature will not, within limits, condensate out liquid particles in the vapor. Highly superheated vapors are gases that approximately follow the general gas law.

7. Critical Temp & Pressure Critical Temperature – The temperature at which only gas exists, regardless of its pressure Critical Pressure – The lowest pressure at which liquids exist at critical temperature Critical Point – The intersection of critical temperature & pressure

8. Kinetic-Molecular Theory of Gases Ideal gas = hypothetical gas perfectly aligns with all kinetic-molecular theory assumptions Five Assumptions – Distance between molecules dwarfs actual size – All collisions are perfectly elastic – Particles are in continuous, rapid, random motion – Particles have NO attraction to each other – Temperature = average kinetic energy of particles

9. Nature of Gases Ideal vs. Real – Real approaches ideal @ low pressure/ high temp Expansion – molecules fill entire space Fluidity – no intermolecular attractions Density - ~ 10 -3 of liquid or solid state Compressibility – 100X more molecules Diffusion & Effusion – Spontaneous mixing via random motion – Passing through tiny opening

10. C/C Intermolecular Forces

11. Properties of Liquids LEAST common state of matter in universe Fluids (as are gases) Lower kinetic energy than gases Interactive forces keep molecules connected – Dipole-dipole forces Equal but opposite charges separated by short distance – London dispersion forces Spontaneous creation of dipoles (polar & nonpolar) – Hydrogen bonding (electronegativity)

12. Hydrogen Bonding

13. Properties of Liquids, continued Density: 100x > gases; 10% < solids Compressibility: @ 10 3 atm., volume ~ 4% Diffusion: present, but slower than in gases Surface tension: high intermolecular attraction Capillary action: attraction between surfaces of liquid and a solid Vaporization: evaporation & boiling  gas

14. Nature of Solids Interparticle attractions stronger than others Two types of solids – Crystalline (orderly arrangement) – Amorphous (random arrangement) supercooled liquids: have liquid properties even if look solid Shape & Volume: Definite Melting Point: Definite Density & Incompressibility: High Diffusion: Low rate (10 -6 less than others)

15. Dipole-dipole Forces

16. Covalent Molecular Structures Buckminsterfullerene Glucose – C 6 H 12 O 6

17. Crystalline Solids Ionic Alkali & alkaline earth with halogens & Group 16 Hard, brittle, high melting points, good insulators Covalent network C x (diamonds), (SiO 2 ) x quartz, (SiC) x Very hard and brittle, high MP, semi- or nonconductors Covalent molecular (nonpolar & polar) H 2, CH 4, C 6 H 6 : only weak London dispersion forces H 2 O & NH 3,: stronger forces but weaker than covalent Soft, low MP, low BP, good insulators

18. Crystalline Structures AgCl Ionic Structure

19. Covalent Network Crystals DiamondQuartz

20. Covalent Network GraphiteCarbon

21. Covalent Molecular