Phases of Matter

1. Phases of Matter

2. Solid • Liquid • Gas • Plasma Phases

3. IMF’s explain the phases of matter • Solid • Particles move relatively slow (basically vibrating in fixed position) • IMF’s hold them together • Liquid • Particles move more (able to flow past each other) • IMF’s still hold together but not as tightly • Gas • Particle move quickly (each in rapid, random motion) • Little to no IMF’s (we usually assume none) Phases of Matter

4. Intermolecular forces are very important. Intermolecular forces are of little significance Intermolecular forces must be considered.

5. Particles relatively close together • Retains volume and shape • Relatively high densities • Hard to compress • Does not flow easily • Types • Crystalline- regular order/ pattern to particles • Amorphous- no regular pattern to particles Solid

6. Particles more spaced apart • Retains volume but not shape • Intermediate densities • Hard to compress • Flows easily (Fluidity) • Slow flow faster than others • Viscosity- measure of the resistance to flow in a liquid • Stronger IMF’s lead to higher viscosity • Can diffuse- liquid molecules spread out through another liquid • Can display surface tension- attraction of molecules at surface to each other (Stronger IMF’s lead to more surface tension) • Can display capillary action- attraction of liquid to solid surface causing it to flow Liquid

8. Particles spaced very far apart • Does not retain volume or shape • Very low densities • Compressible • Flows easily (fluidity) • Can diffuse and go through effusion • Elastic collisions (no loss of energy during collision) • Ideal vs Real Gases • Ideal- no attraction between particles • Real- gases where particles are attracted to each other Gas

9. Occurs at very high temperature and pressures • Also a fluid • Mix of neutral atoms, free electrons, and ions Plasma

10. Density • Physical, intensive property • Mass per unit volume of substance (D= m/V) • Units • Solids and Liquids • g/cm3 or g/ml (1cm3 = 1 ml) • Gases • g/L • Density controls placement of fluids and solids • Less dense objects or fluids move to the top • More dense objects or fluids move to the bottom

11. D = m/V • Mass • Measured on a balance • Volume • Solid • Regular Shape- can be calculated from other measurements • Irregular Shape- can be found by water displacement method • Liquid and Gas • Can be measured with instruments such as graduated cylinder Density (cont)

12. Amount of energy transferred from one substance to another • Represented by q with units in Joules (J) • When heat transfers, it affects the temperatures of the substances Heat

13. Measure of the average kinetic energy in a sample • High temperatures mean the particles are moving quickly • Theoretically if the particles weren’t moving at all, the temperature would be 0 Kelvin (absolute zero) • Remember K = °C + 273.15 Temperature

14. Heat will “flow” from the substance with a higher temperature to the substance with a lower temperature • As the faster particles collide with slower particles, the faster ones will slow down and the slower ones will speed up Heat Flow

15. When heat transfers, it affects the temperatures of the substances involved in the transfer • How much will the temperature change? • Dependent on • Amount of heat transferred • Mass of the sample • Composition of the sample Heat and Temperature Change

16. The more heat transferred, the greater the temperature change • If heat is absorbed by the sample • q is positive • Final temperature will be higher than the initial temperature • If heat is lost by the sample • q is negative • Final temperature will be lower than the initial temperature Amount of Heat Transferred

17. A heat transfer will cause a bigger temperature change to a smaller mass than it will to a larger mass. Mass of Sample

18. Different substances absorb/release heat in different ways. • Specific heat (c) – is the amount of heat needed to change 1 gram of a particular substance by 1 °C. • Each type of substance has a different value Composition of the Sample

19. q= mcT • q is heat • m is mass • c is specific heat • T is change in temperature • Tf-Ti (final temperature – initial temperature) • Units need to match • For instance, if the specific heat value is given in J/g°C, heat should be in J, mass in grams, and temperature in °C Specific Heat Equation

20. Instrument used to measure heat transfer • Process allows water and another substance to undergo a heat transfer until both are at the same temperature (thermal equilibrium) • This means the heat gained (or lost) by the water equals the heat lost (or gained) by the substance Calorimeters

21. Solid absorbs heat and temperature increases (molecules moving faster) • Reaches a point that movement weakens IMF’s enough to allow flow (melting point) • Heat is still absorbed but temperature does not increase • Liquid absorbs heat and temperature increases (molecules moving faster) • Reaches a point that movement weakens IMF’s enough that they essentially no longer exist (boiling point) • Heat is still absorbed but temperature does not increase • Gas absorbs heat and temperature increases (molecules moving faster) Absorbing Heat

22. Plateaus at melting point • Heat still added • Used to weaken IMF’s • Called Heat of fusion • Another longer plateau will occur between liquid and gas • Called heat of vaporization Heating Curve

23. Heat is released from the gas and temperature decreases (molecules moving slower) • Reaches a point that molecules are close enough for IMF’s to be reestablished (condensation point) • Heat is still released but temperature does not decrease • Heat is released from the liquid and temperature decreases (molecules moving slower) • Reaches a point that molecules are close enough for IMF’s to strengthen (freezing point) • Heat is still released but temperature does not decrease • Heat is removed from the solid Releasing Heat

24. Heating/Cooling Curve

25. Solid  Liquid • Solid  Liquid = Melting • Liquid  Solid = Freezing • Occurs at melting/ freezing point • Liquid  Gas • Liquid  Gas = Evaporation/Vaporization/Boiling • Gas  Liquid = Condensation • Occurs at boiling/condensation point • Solid  Gas • Solid  Gas = Sublimation • Gas  Solid = Deposition Phase Changes

26. Evaporation • Occurs at temperatures below boiling point • Some molecules have enough energy to escape surface of liquid • Vaporization • Occurs at boiling point • Change to gaseous phase occurs throughout liquid Evaporation or Vaporization

27. Freezing point and Melting point • Same thing (occur at same temperature) • Named depending on the direction compound is going • Condensation point and Boiling point • Same thing (occur at same temperature) • Named depending on the direction compound is going • Each substance has its own points and heats (fusion and vaporization) Phase Change (cont)

28. Same substance • Pressure (mostly sways boiling point) • “Normal” points are points at standard pressure (1atm) • Lower pressures allow particles to spread out more (IMF’s can be overcome at lower temps) • Higher pressures compress molecules (Higher temp needed to overcome IMF’s) • Between different compounds • Strength of forces holding particles together • Metallic Bonds • Ionic Bonds • Covalent Bonds • IMF’s Influencing Points

30. Vapor pressure- Partial pressure of gas particles of substance over the liquid of that substance • Vapor pressure increases with temperature • More particles have energy to escape surface • Point at which vapor pressure of substance is equal to atmospheric pressure Boiling Point

31. Chart for each substance showing the temperature for phase changes according to pressure • Crossing a line indicates a phase change Phase Diagram

33. At any pressure, a horizontal line can be drawn. • Temperatures of phase changes are found where lines are crossed • The “normal” points are found by drawing a horizontal line at 1atm of pressure Phase Diagrams (cont)

35. Triple Point • Pressure and Temperature where all three phases can be found • Critical Point • Critical Temperature- highest temperature that the liquid phase of a substance can be found • Critical Pressure- pressure at critical point • Beyond this point the liquid and gas phase in indistinguishable (super critical fluid) Phase Diagrams (cont)