1. States of Matter

2. I. Review: Phases of Matter A.Solid –Fixed volume and shape –Molecules are tightly packed and in a set position B. Liquid –Fixed volume, indefinite shape –Molecules are tightly packed, but not in a set position. C. Gas –No fixed volume or shape –Molecules are widely dispersed.

3. II. Forces of Attraction A.Intramolecular Forces –Forces that act within a molecule. Ex: –Covalent and ionic bonds

4. II. Forces of Attraction B. Intermolecular Forces Short range forces between molecules

5. II. Forces of Attraction 1. Hydrogen Bonding A particularly Strong force Occurs in compounds in which a hydrogen atom is attached to Fluorine, Oxygen, or Nitrogen atom. The hydrogen atom will act like a positive ion The hydrogen atom will then be attracted to negative charge on a nearby molecule.

6. II. Forces of Attraction 1. Hydrogen Bonding Hydrogen bonding increases melting and boiling points because more energy is required to break the forces between molecules.

7. II. Forces of Attraction 2. Dipole-Dipole Forces An intermolecular force Occurs in polar molecules (dipoles) because they have uneven charge distribution (a positive end and a negative end) A positive end of one molecule is attracted to a negative end of a nearby molecule. Dipole – dipole forces will also increase melting and boiling points. A dipole can also temporarily attract electrons from another molecule causing an temporary dipole.

8. II. Forces of Attraction 3. London Dispersion Forces A weaker intermolecular force (a type of induced dipole). Occurs in nonpolar molecules and noble gases. Electrons are constantly moving, therefore electron density changes. This effect increases with increasing number of electrons. F2 Gas Br2Liquid I2Solid

9. III. Liquids and Solids A. Liquids 1. Density and Compression a. density similar to that of solids b. Not very compressible 2. Fluidity a. Weak intermolecular forces allow molecules to flow past each other. 3. Viscosity a. The resistance a material has to flowing. b. Increases with increasing intermolecular forces.

10. III. Liquids and Solids B. Solids 1. Density of Solids a. Usually higher than liquids and solids for the same substance. 2. Crystalline Solids a. Solids consisting of ionically bound ions. 3. Network Covalent Solids a. Solids consisting of covalently bound atoms. 4. Metallic Solids a. Metal atoms bound with metallic bonds. 5. Amorphous Solid a. A solid with no molecular pattern.

11. Crystalline solid

12. Network Covalent solids

13. Metallic Solid

14. Amorphous solid

15. IV. Gases and the Kinetic Molecular Theory A.Kinetic Molecular Theory –Many of the properties of gases can be explained by the KMT B. Assumptions: 1.Particle Size Atoms of gases are so much smaller than the spaces between atoms, that their size is negligible. 2. Particle Motion a. Collisions elastic – No energy is lost during a collision. inelastic – some energy is lost during a collision.

16. IV. Gases and the Kinetic Molecular Theory b. Constant, rapid, random motion c. No forces of attraction or repulsion between gas particles

17. IV. Gases and the Kinetic Molecular Theory 3. Particle Energy –Kinetic Energy = (1/2)(Mass) (velocity) 2 –Depends on upon two factors Particle velocity Particle mass

18. IV. Gases and the Kinetic Molecular Theory Temperature 1.Definition: The average kinetic energy of the particle of a substance. 2.Three scales: a. Celsius Based on the boiling and freezing points of water. b. Kelvin Sets to coldest theoretical temperature to “zero.” The difference between two temperatures is the same in both Celsius and Kelvin c. Fahrenheit Based on the human body temperature and the coldest temperature that the creator could get water to freeze at.

19. IV. Gases and the Kinetic Molecular Theory 3. Conversions T f = ((9/5)T c ) + 32 T c = (5/9)(T f – 32) T k = T c + 273

20. IV. Gases and the Kinetic Molecular Theory 4. Practice: a. 28 o C = _____________K b. 200 K = _____________ o C c. – 15 o C = ____________ o F d. 10 K = _______________ o F

21. IV. Gases and the Kinetic Molecular Theory Physical Properties 1.Density – Very Low 2.Fluidity – Flows very easy 3.Compression – Highly compressible 4.Expansion – Expands to fill container 5.Diffusion – The movement from an area of higher concentration to lower concentration. 6.Effusion – The movement from an area of higher concentration to lower concentration through an opening.

22. IV. Gases and the Kinetic Molecular Theory E. Graham’s Law of Diffusion/Effusion Essentially, small molecules move faster than larger ones.

23. IV. Gases and the Kinetic Molecular Theory Sample: Which gas will diffuse the fastest: ammonia, NH3, or hydrogen chloride, HCl?

24. IV. Gases and the Kinetic Molecular Theory F. Gas Pressure 1. Definition of Pressure: The force of molecules on their container 2. Mathematical Formula: P = (Force/Area) 3. Air Pressure Conversions-Conversion Factors 760 mmHg (torr) = 1 atm = 101.3 kPa = 30.0 in Hg = 1 torr =.133 kPa

25. IV. Gases and the Kinetic Molecular Theory Sample: Convert 0.70 atm to mmHg. Practice: Convert 2.0 atm to torr (mmHg). Convert 725 torr to atm.

26. IV. Gases and the Kinetic Molecular Theory 4. Atmospheric (air) Pressure –The pressure that the air around us places on everything 5. Measuring Atmospheric Pressure a. Hg Barometer = glass tube sealed at one end with the other end immersed in a container of Hg

27. A typical Mercury Barometer

28. IV. Gases and the Kinetic Molecular Theory 5. Measuring Pressure of Gases a. Closed Manometer b. Open Manometer

30. IV. Gases and the Kinetic Molecular Theory 5. Standard Pressure 6. Air Pressure Varies with Altitude Dalton’s Law of Partial Pressures Ptotal = P1 + P2 + P3 …….

31. IV. Gases and the Kinetic Molecular Theory Practice: 1. Determine the total pressure, in mm Hg, for a mixture that contains four gases with partial pressures of 5.0 atm, 4.56 atm, 3.02 atm, and 1.2 atm.

32. IV. Gases and the Kinetic Molecular Theory 2. What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600 mmHg and the partial pressure of helium is 439 mmHg?

33. IV. Gases and the Kinetic Molecular Theory V. Phase Changes Phase changes that require energy (endothermic) – Melting – Vaporization – Sublimation

34. IV. Gases and the Kinetic Molecular Theory Phase changes that release energy (exothermic) – Condensation – Deposition – Freezing

35. IV. Gases and the Kinetic Molecular Theory Heating Curve