Phases of matter

1. Phases of matter Energy and phases

2. Everything that has mass and volume is called matter. What is matter?

3. What kind of changes does matter undergo? All matter can undergo changes in its state. These changes have to do with the amount of energy in the particles of matter.

4. Kinetic theory • Kinetic theory of matter- • 1.All matter is made of particles • 2.These particles are in constant motion. • More energy causes the particles to move faster.

5. Kinetic Molecular Model of Water Between 0°C and 100 °C, water is a liquid. In the liquid state, water molecules are close together, but can move about freely. At 100°C, water becomes water vapor, a gas. Molecules can move randomly over large distances. Below 0°C, water solidifies to become ice. In the solid state, water molecules are held together in a rigid structure.

6. Matter: takes up space and has mass • Matter has five states or phases • Solid : A definite shape and volume • Lower energy

7. Liquid: A definite volume but it takes the shape of its container • Higher energy • Gas : No definite shape or volume • Yet even higher energy

8. Special phases • Plasma : No definite shape or volume and whose particles have broken apart • Bose-Einstein Condensate: Gases near absolute zero forming a super fluid

9. Plasma Plasma is by far the most common form of matter. Plasma in the stars and in the tenuous space between them makes up over 99% of the visible universe and perhaps most of that which is not visible.

10. Liquids: Are not very compressible • Useful in hydraulics • Viscosity: Liquids resistance to flow • Surface tension: holds the liquid together

11. Gases: Fill their container and can change pressure • http://www.stolaf.edu/people/giannini/flashanimat/transport/project.swf • Diffusion: Spreading of particles through an area until they are uniformly distributed

12. Changes of state • Changes of state: When matter changes form, it is a physical change and has to do with the energy of the material • Ex. Boiling, melting, freezing, condensing –all require an energy change

13. Melting • Heat of fusion: Energy required to go from solid state to liquid state (For water 334kJ/kg) • Melting point – different for substances

14. Boiling • helium−269 • hydrogen−253 • Iron-2887 • Graphite (carbon)3900 • Diamond (carbon)4827 • Tungsten-5660 • Gold-3080 • Heat of Vaporization: Energy required to go from liquid to gas. (For water 2260kJ/kg)

15. Condensation: A gas changes to a liquid when cooled to or below its boiling point • Vaporization is at boiling point or below • Evaporation: A liquid changes to a gas without reaching its boiling point

16. Skip the liquid • Sublimation: Changing from a solid to a gas without existing as a liquid • Deposition : changing from a gas to a solid without being a liquid

18. Phase change energy • Latent heat: heat absorbed without a change in temperature (stored until a phase change) • Is water a solid or liquid at 0C? Why?

19. Honor only Intermolecular Forces

20. A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. 2 Phases Solid phase - ice Liquid phase - water 11.1

21. Generally, intermolecular forces are much weaker than intramolecular forces. Intermolecular forces are attractive forces between molecules. Intramolecular forces hold atoms together in a molecule. • Intermolecular vs Intramolecular • 41 kJ to vaporize 1 mole of water (inter) • 930 kJ to break all O-H bonds in 1 mole of water (intra) “Measure” of intermolecular force boiling point melting point DHvap DHfus DHsub 11.2

22. Intermolecular forces are feeble; but without them, life as we know it would be impossible. Water would not condense from vapor into solid or liquid forms if its molecules didn't attract each other. Intermolecular forces are responsible for many properties of molecular compounds, including crystal structures (e. g. the shapes of snowflakes), melting points, boiling points, heats of fusion and vaporization, surface tension, and densities. Intermolecular forces pin gigantic molecules like enzymes, proteins, and DNA into the shapes required for biological activity. http://www.nationmaster.com/encyclopedia/Image:Myoglobin.png

23. Intermolecular Forces • 1. London Forces (Dispersion Forces) • Dipole-Dipole Interactions • 3. Ion-Dipole Interactions (Salt dissolving in solution) • 4. Hydrogen Bonding

24. Dispersion Forces Occur between every compound and arise from the net attractive forces amount molecules which is produced from induced charge imbalances Figure 10-8 Olmsted Williams The magnitude of the Dispersion Forces is dependent upon how easily it is to distort the electron cloud. The larger the molecule the greater it’s Dispersion Forces are. Figure 10-9 Olmsted Williams

25. The boiling point of alkanes increase with the length of the carbon chain. Long-chain alkanes have larger dispersion forces because of the increased polarizability of their larger electron cloud. Olmsted Williams Fig 10-10 Pg 437

26. How molecular shape affects the strength of the dispersion forces The shapes of the molecules also matter. Long thin molecules can develop bigger temporary dipoles due to electron movement than short fat ones containing the same numbers of electrons. Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close. For example, the hydrocarbon molecules butane and 2-methylpropane both have a molecular formula C4H10, but the atoms are arranged differently. In butane the carbon atoms are arranged in a single chain, but 2-methylpropane is a shorter chain with a branch. Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules. http://www.chemguide.co.uk/atoms/bonding/vdw.html

27. Dispersion forces usually increase with molar mass. Polarizability the ease with which the electron distribution in the atom or molecule can be distorted. • Polarizability increases with: • greater number of electrons • more diffuse electron cloud 11.2

28. Is the Molecule Polar? We have already talked about diatomic molecules. The more Electronegative atom will pull the electron density of the bond Closer to itself giving it a partial negative charge leaving the other Atom with a partially positive charge. Thus giving the molecule A dipole moment. But what about molecules made up of more than two molecules?

29. Orientation of Polar Molecules in a Solid Dipole-Dipole Forces Attractive forces between polar molecules 11.2

30. Dipole Forces occur between molecules containing a dipole moment. The positive end of the dipole moment on one mole is attracted to the Negative end of the dipole moment on a nearby molecule. Consider 2-methyl propane (left) and acetone (right) Both compounds are about Equal in size and shape therby Having similar dispersion forces, But Acetone contains an Oxygen (red) and causes the Molecule to have a dipole Moment allowing it to have Dipole forces and thus a Higher boiling point Figure 10-11 Olmsted Williams

31. Ion-Dipole Interaction Ion-Dipole Forces Attractive forces between an ion and a polar molecule The larger the charge the stronger the force 11.2

32. Fig 10-34 Olmsted Williams A molecular picture showing the ion-dipole Interaction that helps a solid ionic crystal dissolve in water. The arrows indicate ion-dipole interactions.

33. O O S What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: dispersion forces. SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules. 11.2

34. or … … H H B A A A Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A & B are N, O, or F 11.2

35. Intermolecular Forces • 1. London Forces (Dispersion Forces) • Dipole-Dipole Interactions • 3. Ion-Dipole Interactions (Salt dissolving in solution) • 4. Hydrogen Bonding These forces affect how molecules will interact with each other and As a general rule as the strength of the force increases the boiling Point of the compound increases

36. Liquids and Surface Tension Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Strong intermolecular forces High surface tension 11.3

37. Adhesion Cohesion Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules 11.3

38. Least Order Greatest Order T2 > T1 Condensation Evaporation 11.8

39. H2O (l) H2O (g) Dynamic Equilibrium Rate of evaporation Rate of condensation = The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation A substance with a high Vapor pressure is considered To be volitile therefore, the lower The boiling point the higher the Vapor pressure and the weaker The intermolecular forces 11.8

40. The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm. 11.8

41. H2O (s) H2O (l) The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium Freezing Melting 11.8

42. 11.8

43. H2O (s) H2O (g) Molar heat of sublimation (DHsub) is the energy required to sublime 1 mole of a solid. Sublimation Deposition DHsub = DHfus + DHvap ( Hess’s Law) 11.8

44. A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas. Phase Diagram of Water 11.9

45. 11.9