Chapter 12

1. Chapter 12 States of Matter

2. 12.1 Gasses • Gasses expand, diffuse, exert pressure, and can be compressed because they are in a low density state consisting of tiny, constantly moving particles

3. Kinetic Molecular Theory • Describes the behavior of matter in terms of particles in motion • Makes several assumptions about the size, motion, and energy of gas particles

4. Assumptions of the Kinetic Molecular Theory • Gasses consist of small particles • Gasses take up little volume relative to the volume of empty space around them so size of individual particles is zero

5. Gas particles move in constant, random straight lines until they collide with other particles or with the walls of the container • Collisions are elastic – • Collisions cause gas pressure -

6. Gas particles do not attract or repel each other • The average kinetic energy of gas particles is directly proportional to the KELVIN temperature of the gas

7. Explaining the Behavior of Gasses • Kinetic molecular theory helps explain the behavior of gasses • Blowing up a balloon

8. Gasses have low densities • Why?

9. Gasses can be compressed and can expand • Explain

10. Gasses diffuse and effuse • Diffusion = gas particles move from an area of high concentration to low concentration • Ex • Effusion = gas particles escape through tiny openings • Ex

11. 12.2 Forces of attraction • Intermolecular forces (dispersion forces, dipole-dipole forces, and hydrogen bonds) determine a substance’s state at a given temperature

12. Intermolecular forces • Inter- means between or among • Intermolecular forces can hold together identical particles or two different types of particles • Weaker than intramolecular forces (bonds)

13. Dispersion Forces • Weak forces that result from temporary shifts in the density of electrons in electron clouds

14. Exist between all particles • Weak for small particles • Get stronger as the number of electrons involved increases • F2 • Cl2 • Br2 • I2

15. Dipole-dipole forces • Attraction between oppositely charged regions of polar molecules • Polar molecule = • Neighboring polar molecules orient themselves so that oppositely charged regions align

17. Hydrogen Bonds • Dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a flourine, oxygen, or nitrogen atom

18. Explain why water is liquid at room temperature while compounds of similar masses are gasses

19. 12.3 Liquids and Solids • The particles in solids and liquids have a limited range of motion and are not easily compressed.

20. Liquids • Kinetic molecular theory also applies to liquids and solids • Must take intermolecular forces into account to apply it

21. Density and compression • Much denser than gasses • Due to intermolecular forces holding particles together • Incompressible • Why can you compress a gas but not a liquid?

22. Fluidity – both gases and liquids are classified as fluids because they can flow and diffuse • Liquids diffuse more slowly because intermolecular attractions interfere with the flow

23. Viscosity - measure of the resistance of a liquid to flow • Attractive forces – stronger intermolecular forces = higher viscosity • Particle size – larger molecules = higher viscosity • Temperature – lower temperature = higher viscosity

24. Surface tension – the energy required to increase the surface area of a liquid by a given amount • Caused by intermolecular forces pulling down on the particles on the surface of a liquid which stretches it tight like a drum

25. Stronger the attraction between particles in a liquid = greater surface tension • Surfactant – lowers the surface tension of water by disrupting hydrogen bonds between water molecules

26. Cohesion – force of attraction between identical molecules • Adhesion – force of attraction between molecules that are different

27. Solids • Solid particles have as much kinetic energy as liquids or gasses but much stronger attractive forces between particles • Limit the motion of particles to vibrations

28. Density of solids – almost always greater than density of liquids Exception = water

29. Crystalline solids – solid whose atoms, ions, or molecules are arranged in an orderly, geometric structure • Unit cell = smallest arrangement of atoms in a crystalline solid that has the same shape as the whole crystal

31. Categories of crystalline solids • Classified based on the types of particles they contain and how they are bonded together

32. Molecular solids • Molecules are held together by dispersion forces, dipole-dipole forces or hydrogen bonds • Most are not solid at room temperature • Poor conductors

33. Covalent network solids • C or Si, can form multiple covalent bonds which allow it to take many forms • Allotrope – element that can exist in different forms at the same state

34. Ionic solids • Made of cation + anion • Each ion is surrounded by ions of the opposite charge • High melting point • Brittle

35. Metallic solids • Positive metal ions surrounded by a sea of mobile electrons

37. Amorphous solids • Particles are not arranged in a regular, repeating pattern • Does not contain crystals • Forms when molten material cools too quickly for crystals to form • Glass • Rubber • Some plastics

38. 12.4 Phase changes • Matter changes phases when energy is added or removed

39. Phase changes that require energy • Melting • Heat flows from an object at a higher temperature to an object at a lower temperature • Ice absorbs heat which does not raise temperature but is used to break hydrogen bonds • When hydrogen bonds are broken molecules can move further apart into the liquid phase

40. Melting point – temperature in which forces holding a solid together are broken and it becomes a liquid

41. Vaporization – process by which liquid changes to vapor • Vapor – gaseous state of a substance that is normally liquid at room temperature • Evaporation – when vaporization occurs only at the surface of a liquid • Vapor pressure – the pressure exerted by a vapor over a liquid

42. Boiling – temperature at which the vapor pressure of a liquid equals the atmospheric pressure • Energy being input causes molecules to move around more and vaporize

43. Sublimation – changing from solid to gas without becoming a liquid • Dry ice • Moth balls • Solid air fresheners

44. Heating Curve

45. Phase changes that release energy • Freezing • Heat flows out of warmer object into cooler object • Molecules slow down & become less likely to flow past one another • Intermolecular forces cause the molecules to become fixed into set positions • Freezing point – temperature in which a liquid becomes a solid

46. Condensation – process by which a gas or vapor becomes a liquid • Deposition – substance changes from gas or vapor to solid without first becoming a liquid • frost

47. Phase Diagrams • Temperature and pressure both effect the phase of a substance • Have opposite effects • Phase diagram – graph of pressure vs temperature that shows which phase a substance will be in under different conditions.

48. Triple point = point at which all three phases exist at the same time