Chapter 10

1. Chapter 10 Intermolecular forces and phases of matter Why does matter exist in different phases? What if there were no intermolecular forces? The ideal gas

2. Physical phases of matter • Gas • Liquid • Solid • Plasma

3. Physical properties of the states of matter Gases: Highly compressible Low density Fill container completely Assume shape of container Rapid diffusion High expansion on heating

4. Liquid (condensed phase) Slightly compressible High density Definite volume, does not expand to fill container Assumes shape of container Slow diffusion Low expansion on heating

5. Solid (condensed phase) Slightly compressible High density Rigidly retains its volume Retains its own shape Extremely slow diffusion; occurs only at surfaces Low expansion on heating

6. Why water exists in three phases? • Kinetic energy(the state of substance at room temperature depends on the strength of attraction between its particles) • Intermolecular forces stick molecules together (heating and cooling)

7. Intermolecular forces • London Force or dispersion forces • Dipole-dipole • Hydrogen bond

8. London Force Weak intermolecular force exerted by molecules on each other, caused by constantly shifting electron imbalances. This forces exist between all molecules. Polar molecules experience both dipolar and London forces. Nonpolar molecules experience only London intermolecular forces

9. Dipole-dipole • Intermolecular force exerted by polar molecules on each other. • The name comes from the fact that a polar molecule is like an electrical dipole, with a + charge at one end and a - charge at the other end. The attraction between two polar molecules is thus a "dipole-dipole" attraction.

10. Hydrogen bond • Intermolecular dipole-dipole attraction between partially positive H atom covalently bonded to either an O, N, or F atom in one molecule and an O, N, or F atom in another molecule.

11. To form hydrogen bonds, molecules must have at least one of these covalent bonds: • H-N or H-N= • H-O- • H-F

14. Nonmolecular substances • Solids that don’t consist of individual molecules. • Ionic compounds(lattices of ions) • They are held together by strong ionic bonds • Melting points are high

15. Other compounds • Silicon dioxide(quartz sand) and diamond (allotrope of carbon) • These are not ionic and do not contain molecules • They are network solids or network covalent substances

16. Real Gas • Molecules travel fast • Molecules are far apart • Overcome weak attractive forces

17. Ideal Gas • Gas that consists of particles that do not attract or repel each other. • In ideal gases the molecules experience no intermolecular forces. • Particles move in straight paths. • Does not condense to a liquid or solid.

18. Kinetic Molecular Theory • Particles in an ideal gas… • have no volume. • have elastic collisions. • are in constant, random, straight-line motion. • don’t attract or repel each other. • have an avg. KE directly related to Kelvin temperature.

19. Kinetic Molecular Theory Postulates of the Kinetic Molecular Theory of Gases • Gases consist of tiny particles (atoms or molecules) • These particles are so small, compared with the distances between • them, that the volume (size) of the individual particles can be assumed • to be negligible (zero). • 3. The particles are in constant random motion, colliding with the walls of • the container. These collisions with the walls cause the pressure exerted • by the gas. • 4. The particles are assumed not to attract or to repel each other. • 5. The average kinetic energy of the gas particles is directly proportional • to the Kelvin temperature of the gas

20. AS TEMP. , KE Kinetic Molecular Theory (KMT) • explains why gases behave as they do • deals w/“ideal” gas particles… 1. …are so small that they are assumed to have zero volume • …are in constant, straight-line motion • …experience elastic collisionsin which no energy is lost • …have no attractive or repulsiveforces toward each other • …have an average kinetic energy (KE)that is proportional • to theabsolute temp. of gas (i.e., Kelvin temp.)

21. Elastic vs. Inelastic Collisions 8 3

22. 8 8 Elastic vs. Inelastic Collisions POW v1 v2 elastic collision v3 v4 inelastic collision

23. 8 8 Elastic Collision v1 before v2 after

24. Ideal Gases Ideal gases are imaginary gases that perfectly fit all of the assumptions of the kinetic molecular theory. • Gases consist of tiny particles that are far apart • relative to their size. • Collisions between gas particles and between • particles and the walls of the container are • elastic collisions • No kinetic energy is lost in elastic • collisions

25. Ideal Gases (continued) • Gas particles are in constant, rapid motion. They • therefore possess kinetic energy, the energy of • motion • There are no forces of attraction between gas • particles • The average kinetic energy of gas particles • depends on temperature, not on the identity • of the particle.

26. Measurable properties used to describe a gas: • Pressure (P) P=F/A • Volume (V) • Temperature (T) in Kelvins • Amount (n) specified in moles

27. Pressure • Is caused by the collisions of molecules with the walls of a container • is equal to force/unit area • SI units = Newton/meter2 = 1 Pascal (Pa) • 1 standard atmosphere = 101.3 kPa • 1 standard atmosphere = 1 atm = • 760 mm Hg = 760 torr

28. Measuring Pressure The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17th century. The device was called a “barometer” • Baro= weight • Meter= measure

29. An Early Barometer The normal pressure due to the atmosphere at sea level can support a column of mercury that is 760 mm high.

30. U-tube Manometer • Manometer • measures contained gas pressure

31. AIR PRESSURE Hg HEIGHT DIFFERENCE manometer: measures the pressure of a confined gas CONFINED GAS SMALL + HEIGHT = BIG differential manometer manometers can be filled with any of various liquids

32. Pressure • Atmosphere Pressure and the Barometer • The pressures of gases not open to the atmosphere are measured in manometers. • A manometer consists of a bulb of gas attached to a U-tube containing Hg: • If Pgas < Patm then Pgas + Ph2 = Patm. • If Pgas > Patm then Pgas = Patm + Ph2.

33. The Gas Laws • The Pressure-Volume Relationship: Boyle’s Law • Weather balloons are used as a practical consequence to the relationship between pressure and volume of a gas. • As the weather balloon ascends, the volume decreases. • As the weather balloon gets further from the earth’s surface, the atmospheric pressure decreases. • Boyle’s Law: the volume of a fixed quantity of gas is inversely proportional to its pressure. • Boyle used a manometer to carry out the experiment. Chapter 10

35. BIG BIG = small + height 760 mm Hg 112.8 kPa 846 mm Hg = height = BIG - small 101.3 kPa X mm Hg 846 mm Hg 593 mm Hg - = X mm Hg = 253 mm Hg 253 mm Hg STEP 1) Decide which pressure is BIGGER STEP 2) Convert ALL numbers to the unit of unknown STEP 3) Use formula Big = small + height small 0.78 atm height X mm Hg 760 mm Hg 0.78 atm 593 mm Hg = 1 atm

36. 96.5 kPa X atm 233 mm Hg Atmospheric pressure is 96.5 kPa; mercury height difference is 233 mm. Find confined gas pressure, in atm. S B SMALL + HEIGHT = BIG + = 96.5 kPa 233 mm Hg X atm 0.953 atm + 0.307 atm = 1.26 atm

37. Units of Pressure

38. Standard Temperature and Pressure“STP” • P = 1 atmosphere, 760 torr, 101.3 kPa • T = 0°C, 273 Kelvins • The molar volume of an ideal gas is 22.4 liters at STP

39. Behavior of gases • Rule 1: P is proportional to 1/V • Rule 2: P is proportional to T • Rule 3: P is proportional to n Combining all three: P is proportional to nT/V P=constant x nT/v R=constant= 0.0821 L atm/K mole

40. PTV PT V P T V 1 V Boyle’s P a ___ a Charles V T Gay-Lussac’s P T a Pressure - Temperature - Volume Relationship

41. Boyle’s Law • P inversely proportional to V • PV= k • Temperature and number of moles constant

42. The Gas Laws • The Pressure-Volume Relationship: Boyle’s Law • Mathematically: • A plot of V versus P is a hyperbola. • Similarly, a plot of V versus 1/P must be a straight line passing through the origin. Chapter 10

43. Boyle’s Law Pressure is inversely proportional to volume when temperature is held constant.

44. A Graph of Boyle’s Law

45. Charles’s Law • V directly proportional to T • T= absolute temperature in kelvins • V/T =k2 • Pressure and number of moles constant

46. The Gas Laws • The Temperature-Volume Relationship: Charles’s Law • We know that hot air balloons expand when they are heated. • Charles’s Law: the volume of a fixed quantity of gas at constant pressure increases as the temperature increases. • Mathematically: Chapter 10

47. Charles’s Law • The volume of a gas is directly proportional to temperature, and extrapolates to zero at zero Kelvin. • (P = constant) Temperature MUST be in KELVINS!

48. A Graph of Charles’ Law

49. The Gas Laws • The Temperature-Volume Relationship: Charles’s Law • A plot of V versus T is a straight line. • When T is measured in C, the intercept on the temperature axis is -273.15C. • We define absolute zero, 0 K = -273.15C. • Note the value of the constant reflects the assumptions: amount of gas and pressure. Chapter 10

50. Gay Lussac’s Law The pressure and temperature of a gas are directly related, provided that the volume remains constant. Temperature MUST be in KELVINS!