Unit 4: Phases of Matter (Chapters 13-14)

1. Unit 4: Phases of Matter(Chapters 13-14)

2. 14.1 Phase Changes • All substances can exist as solids, liquids, and gases. • States of Matter are affected by: • IMF • KMT

3. How does structure of matter determine its natural state? • What elements are solids, liquids, and gases at room temperature? • What compounds are solids, liquids, and gases at room temperature? • Why?

4. 14.3 Intermolecular Forces • INTRAmolecular vs. INTERmolecular Forces • Intermolecular forces are BETWEEN compounds. • Ionic Electrostatic attractions (ionic compounds) • Dipole-dipole attraction (polar covalent compounds) • Hydrogen bonding (between H and NOF) • London dispersion forces (exist in all molecules including nonpolar compounds and noble gases)

5. IMF • Ionic Electrostatic Attractions • Strongest of all intermolecular forces. • Most ionic compounds are found in the solid form.

6. Dipole-dipole interaction: • Weaken as distance increases • Often found as liquids

7. Hydrogen bonding: • Hydrogen is bound to highly electronegative atom • A particularly strong type of dipole-dipole because: • Great polarity of the bond itself • Proximity of dipoles due to small size of hydrogen • Results in the unusually high boiling point of water

8. London Dispersion Forces- • weakest IMF • Electrons are constantly moving and may become unevenly distributed. • Forces become stronger with more electrons • Often found as gases

9. Main Ideas of IMF on States • What is the difference between intermolecular and inter-ionic forces? • Intermolecular forces are typically weak • Exist as liquids or gases at room temperature • Inter-ionic forces are relatively strong • Exist as solids at room temperature

10. How do pressure, temperature, and volume affect states of matter? • Describe how various states of matter move.

11. 13.8 Kinetic Molecular Theory • Particles are so small next to the space between them that their volume can be assumed to be negligible • Particles are in constant motion- collisions exert pressure • Particles assumed to exert no forces on each other • Average KE of collection of particles is assumed to be directly proportional to K temperature of the gas

12. States of Matter and KMT • Based on KMT, how are states of matter affected by: • Temperature • Pressure • Volume

13. Temperature • How does IMF effect boiling and melting point? • Intermolecular - leads to a relatively low boiling and melting point • Inter-ionic - leads to solids with high melting and boiling points. • How does temperature affect the motion of particles?

14. 2.7 Temperature Scales Fahrenheit Scale Water Boils @ Water Freezes @ Celcius Scale Water Boils @ Water Freezes @ Kelvin Scale Water Boils @ Water Freezes @ Absolute zero is

15. 10.2 Changes in State of Matter How are heat and temperature related to changes in states of matter?

16. 14.4 Interconversion of States • Vaporization/evaporation- when molecules of a liquid escape from the liquids surface and form a gas • Condensation- process by which vapor/gas molecules reform a liquid • Sublimation- process which solids change directly to a gas

17. 14.1 Heating Curve • Heating curve- plot of temperature versus time where energy is added at a constant rate Where on the graph is kinetic energy represented, potential energy, and heat transfer?

18. Special Cases • Supercooled- can be cooled below 0°C and remain a liquid • Superheated- raised to temperatures above boiling point

19. Practice Problems (Honors) • Heat required when a substance undergoes any combination of temperature change and phase change (see Example 14.1 and 14.2).

20. Pressure and Volume • How do pressure and volume affect states of matter?

21. 13.1 Pressure – Units of Pressure in Chemistry 1. mm Hg or torr 2. atmospheres (atm) 3. pascals (Pa) 4. Newtons per square meter (N/m2) 1 atm = 760 mm Hg = 760 torr = 101,325 Pa Volume – Units:

22. 14.4 Phase Diagrams • Show phase changes that occur at equilibrium • Triple point • Critical point • Normal melting point • Normal boiling point • Vapor Pressure

23. Figure 10.50 Diagrams of Various Heating Experiments on Samples of Water in a Closed System

24. KMT and IMF • Use KMT and IMF to explain

25. Crystalline solids are composed of highly regular arrangement of their components Amorphous solids have a considerable amount of disorder in their structures 14.5 SOLIDS

27. 14.6 Ionic Solids • Properties • Stable • High melting points due to strong electrostatic forces between ions • If dissolved in water, the solid dissociates and can conduct electricity. • Structure • Smaller cations fit into holds to maximize electrostatic attractions among ions and minimize repulsions

28. 14.6 Molecular Solids • Strong covalent bonding within molecules • Weak forces between molecules • Requires less energy to break the intermolecular forces. • Dipole IMF: low mpnt and bpnt • H Bonding IMF: moderate mpnt and bpnt • London Dispersion: • Small molecules: gas at room temp • Large molecules: gas or liquid at room temp. • If dissolved in water, the solid splits into individual molecules and does not conduct electricity.

29. 14.6 Atomic Solids • In metallic solids, a special type of delocalized non-directional covalent bonding occurs • In network solids, the atoms bond to each other with strong directional covalent bonds that lead to giant molecules, or networks, of atoms • In the Group 8A solids, the noble gas elements are attracted to each other with London dispersion forces

30. 14.6 Metal Atomic Solids • What is the electron sea model? • Metal cations surrounded by a “sea” of electrons • Properties • malleable, good conductors, and ductile • Electrons are free to move around transmitting heat and electricity • Atoms are easily moved in relation to each other but hard to separate

31. Chapter 13 - Gases

32. 13.9 Implications of Kinetic Molecular Theory for Gases A. The meaning of temperature – B. The relationship between pressure and temperature – C. The relationship between volume and temperature –

33. Figure 13.3: A device (called a manometer) for measuring the pressure of a gas in a container.

34. Figure 13.1: The pressure exerted by the gases in theatmosphere can be demonstrated by boiling water ina can, and then turning off the heat and sealing the can. Hmco Photo Files

35. 13.2 Pressure and Volume (Boyle’s Law) • Pressure and volume are inversely related. • Makes sense because: decreasing volume means that particles will hit the walls of the container more often- causing more pressure • *Boyle’s Law: P1V1 = P2V2

36. Figure 13.6: An illustration of Boyle's Law.

37. The air in a balloon expands when it is heated. Thismeans that some of the air escapes from the balloon, lowering the air density inside and thus making the balloon buoyant. Hmco Photo Library (Royalty Free)

38. Pressure and Temperature • Pressure is directly proportional to temperature • Temperature is a measure of kinetic energy. • Makes sense because: when temperature increases, speed of particles increase, so particles hit wall with greater force, so pressure increases

39. 13.3 Volume and Temperature (Charles’s Law) • Volume is directly proportional to temperature • Makes sense because: Higher temperature = higher speed = hit walls more often with more force= causes expansion of the container the volume is the same • V1=V2 T1 T2

40. Figure 13.13: An increase intemperature results in an increase in pressure.

41. 13.4 Volume and Number of Moles (Avogadro’s Law) • Volume is related directly to the number of gas particles present • Makes sense because: increase in number of particles increases pressure because of increased numbers of collisions • Also volume does not depend on type of particle but on number of particles because volume of individual particles is insignificant

42. 13.5 Ideal Gas Law A. PV = nRT , where R is the universal gas constant (R = 0.08206 L atm/K mol) Note: the units of pressure, volume, and Temperature that you plug into the above equation must match the corresponding units in the R value) p. 402 – Example 13.8 and 13.9 B. The ideal gas law has limitations, It assumes the following:

43. C. It is important to note that the idealgas law is an approximation. Real gases may deviate from the way this model says that they should behave, but the model gets better at extremely low pressures and/or high temperatures. *Self Check exercises 13.5 and 13.6 on p. 403 (Practice – p. 420 45-64)

44. Density and the Ideal Gas Law • You can use the density of a gas to calculate its molar mass.

45. 13.10 Gas Stoichiometry A. Remember…Stoichiometry is all about using a mole ratio. Well…how can you get moles of gas from other information given about the gas? Ans: often you can use PV = nRT to find moles of the gas given before doing a mole ratio. Example 13.15 on p. 414 Self Check Exercise 13.11 on p. 415 Example 13.16 on p 415 Self Check Exercise 13.12 on p. 415 (Practice – p. 422, 83-102)