1. Solid, Liquids, and Gases Their properties and changes

2. Properties of Solids, Liquids, and Gases What are some of the characteristics that distinguish solids, liquids, and gases? What is the evidence for each of these properties?

3. PropertySolidLiquidGas Shape Volume Compressibility Space between particles

4. PropertySolidLiquidGas Relative Density Fluid? Diffusion rate Motion of particles (amount and type) Forces between particles?

5. The three states of matter. Silberberg, Principles of Chemistry

6. STATES OF MATTER SOLIDS Particles of solids are tightly packed, vibrating about a fixed position. Solids have a definite shape and a definite volume. Heat

7. STATES OF MATTER LIQUID  Particles of liquids are tightly packed, but are far enough apart to slide over one another.  Liquids have an indefinite shape and a definite volume. Heat

8. STATES OF MATTER GAS  Particles of gases are very far apart and move freely.  Gases have an indefinite shape and an indefinite volume. Heat

9. PHASE CHANGES Description of Phase Change Term for Phase Change Heat Movement During Phase Change Solid to liquid Melting Heat goes into the solid as it melts. Liquid to solid Freezing Heat leaves the liquid as it freezes.

10. PHASE CHANGES Description of Phase Change Term for Phase Change Heat Movement During Phase Change Liquid to gas Vaporization, which includes boiling and evaporation Heat goes into the liquid as it vaporizes. Gas to liquidCondensation Heat leaves the gas as it condenses. Solid to gasSublimation Heat goes into the solid as it sublimates. Gas to soliddeposition Heat leaves the gas usually at very low pressure

11. STATES OF MATTER PLASMA  A plasma is an ionized gas.  A plasma is a very good conductor of electricity and is affected by magnetic fields.  Plasmas, like gases have an indefinite shape and an indefinite volume. Plasma is the common state of matter

12. STATES OF MATTER SOLID LIQUID GAS PLASMA Tightly packed, in a regular pattern Vibrate, but do not move from place to place Close together with no regular arrangement. Vibrate, move about, and slide past each other Well separated with no regular arrangement. Vibrate and move freely at high speeds Has no definite volume or shape and is composed of electrical charged particles

13. Gases Because gases have so much space between the particles they have properties that are dependent on one another.

14. Gas Variables Volume (V) - mL, L, kL… Temperature (T) – o C measured in lab but K (kelvin) for calculations Number of particles (n) – moles Pressure (P) – mmHg, psi…(more to come)

15. Pressure Force per unit area

16. Figure 5.3 A mercury barometer Silberberg, Principles of Chemistry

17. Table 5.1 Common Units of Pressure Atmospheric PressureUnitScientific Field chemistryatmosphere(atm)1 atm* pascal(Pa); kilopascal(kPa) 1.01325x10 5 Pa; 101.325 kPa SI unit; physics, chemistry millimeters of mercury(Hg) 760 mm Hg*chemistry, medicine, biology torr760 torr*chemistry pounds per square inch (psi or lb/in 2 ) 14.7lb/in 2 engineering bar1.01325 barmeteorology, chemistry, physics *This is an exact quantity; in calculations, we use as many significant figures as necessary. Silberberg, Principles of Chemistry

18. John A. Schreifels Chemistry 211 Chapter 11-18 Phase Diagrams Graph of pressure-temperature relationship; describes when 1,2,3 or more phases are present and/or in equilibrium with each other. Lines indicate equilibrium state two phases. Triple point- Temp. and press. where all three phases co-exist in equilibrium. Critical temp.- Temp. where substance must always be gas, no matter what pressure. Critical pressure- vapor pressure at critical temp. Critical point- point where system is at its critical pressure and temp.

19. Gas Variable Relationships To investigate the relationship between 2 gas variables we need to hold the other 2 constant. Constant P - same # of collisions/unit area Constant V - rigid container Constant T – thermostat control Constant n – keep container sealed

20. The Relationship Between Pressure and Volume Silberberg, Principles of Chemistry

21. The relationship between the volume and pressure of a gas. Boyle’s Law Silberberg, Principles of Chemistry

22. A molecular description the relationship between temperature and volume.

23. The relationship between the volume and temperature of a gas. Charles’s Law

24. The relationship between pressure and temperature As temperature increases, gas particles move faster and make more collisions. As a result the pressure in the container increases. For an aerosol can the pressure may be so great that the seam on the can may give way in an explosion.

25. An experiment to study the relationship between the volume and amount of a gas. The more gas particles you have the more collisions occur. To keep the pressure the same, the volume has to increase so there is more room for the particles. This is why balloons expand when you blow air into them.

26. Solids and Liquids Because the particles are so much closer in liquids and solids, there are chances for particles to attract (or repel). This and the mass of the particles are main factors in determining the properties of solids and liquids. Some properties are boiling and melting points, surface tension, vapor pressure, and crystalline structure.

27. Surface Tension Surface tension is the tendency for liquid surface to contract. Depends on attractive forces Compounds that interfere with the forces and reduce surface tension are called surfactants.

28. The molecular basis of surface tension. hydrogen bonding occurs in three dimensions hydrogen bonding occurs across the surface and below the surface the net vector for attractive forces is downward

29. Shape of water or mercury meniscus in glass. adhesive forces stronger cohesive forces H2OH2O capillarity Hg

30. Solids Solids may have a definite structure and are called crystalline. Solids that have no regular shape are called amorphous.

31. The hexagonal structure of ice.

32. The striking beauty of crystalline solids.

33. portion of a 3-D lattice The crystal lattice and the unit cell. lattice point unit cell portion of a 2-D lattice unit cell

34. Phase Changes

35. solidliquidgas melting freezing vaporizing condensing sublimination Energy absorbed Energy released

36. John A. Schreifels Chemistry 211 Chapter 11-36 Overview Changes of State – Phase transitions – Phase Diagrams Liquid State – Properties of Liquids; Surface tension and viscosity – Intermolecular forces; explaining liquid properties Solid State – Classification of Solids by Type of Attraction between Units – Crystalline solids; crystal lattices and unit cells – Structures of some crystalline solids – Calculations Involving Unit-Cell Dimensions – Determining the Crystal Structure by X-ray Diffraction

37. John A. Schreifels Chemistry 211 Chapter 11-37 Comparison of Gases, Liquids and Solids – Gases are compressible fluids. Their molecules are widely separated. – Liquids are relatively incompressible fluids. Their molecules are more tightly packed. – Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move. Figure 11.2 States of Matter

38. John A. Schreifels Chemistry 211 Chapter 11-38 Phase Transitions Melting: change of a solid to a liquid. Freezing: change a liquid to a solid. Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called sublimation. Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called deposition. H 2 O(s)  H 2 O(l) H 2 O(l)  H 2 O(s) H 2 O(l)  H 2 O(g) or H 2 O(s)  H 2 O(g) H 2 O(g)  H 2 O(l) or H 2 O(g)  H 2 O(s)

39. John A. Schreifels Chemistry 211 Chapter 11-39 Vapor Pressure In a sealed container, some of a liquid evaporates to establish a pressure in the vapor phase. Vapor pressure: partial pressure of the vapor over the liquid measured at equilibrium and at some temperature. Dynamic equilibrium

40. John A. Schreifels Chemistry 211 Chapter 11-40 Temperature Dependence of Vapor Pressures The vapor pressure above the liquid varies exponentially with changes in the temperature. The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

41. John A. Schreifels Chemistry 211 Chapter 11-41 Clausius – Clapeyron Equation A straight line plot results when ln P vs. 1/T is plotted and has a slope of  H vap /R. Clausius – Clapeyron equation is true for any two pairs of points. Write the equation for each and combine to get:

42. John A. Schreifels Chemistry 211 Chapter 11-42 Using the Clausius – Clapeyron Equation Boiling point the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere. Normal boiling point the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm). E.g. Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of 190.0 mmHg at 25.0°C. E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it has a vapor pressure of 100.0 mmHg. What is the heat of vaporization?

43. John A. Schreifels Chemistry 211 Chapter 11-43 Energy of Heat and Phase Change Heat of vaporization: heat needed for the vaporization of a liquid. H 2 O(l)  H 2 O(g)  H = 40.7 kJ Heat of fusion: heat needed for the melting of a solid. H 2 O(s)  H 2 O(l)  H = 6.01 kJ Temperature does not change during the change from one phase to another. E.g. Start with a solution consisting of 50.0 g of H 2 O(s) and 50.0 g of H 2 O(l) at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water.

44. John A. Schreifels Chemistry 211 Chapter 11-44 Properties of Liquids Surface tension: the energy required to increase the surface area of a liquid by a unit amount. Viscosity: a measure of a liquid’s resistance to flow. Surface tension: The net pull toward the interior of the liquid makes the surface tend to as small a surface area as possible and a substance does not penetrate it easily. Viscosity: Related to mobility of a molecule (proportional to the size and types of interactions in the liquid). –Viscosity decreases as the temperature increases since increased temperatures tend to cause increased mobility of the molecule.

45. John A. Schreifels Chemistry 211 Chapter 11-45 Intermolecular Forces Intermolecular forces: attractions and repulsions between molecules that hold them together. Intermolecular forces (van der Waals forces) hold molecules together in liquid and solid phases. – Ion-dipole force: interaction between an ion and partial charges in a polar molecule. – Dipole-dipole force: attractive force between polar molecules with positive end of one molecule is aligned with negative side of other. – London dispersion Forces: interactions between instantaneously formed electric dipoles on neighboring polar or nonpolar molecules. – Polarizability: ease with which electron cloud of some substance can be distorted by presence of some electric field (such as another dipolar substance). Related to size of atom or molecule. Small atoms and molecules less easily polarized.

46. John A. Schreifels Chemistry 211 Chapter 11-46 Boiling Points vs. Molecular Weight Hydrogen bonds: the interaction between hydrogen bound to an electronegative element (N, O, or F) and an electron pair from another electronegative element. Hydrogen bonding is the dominate force holding the two DNA molecules together to form the double helix configuration of DNA.

47. John A. Schreifels Chemistry 211 Chapter 11-47 Comparisonof Energies for Intermolecular Forces Interaction Forces:Approximate Energy Intermolecular London1 – 10 kJ Dipole-dipole3 – 4 kJ Ion-dipole5 – 50 kJ Hydrogen bonding10– 40 kJ Chemical bonding Ionic100 – 1000 kJ Covalent100 – 1000 kJ

48. John A. Schreifels Chemistry 211 Chapter 11-48 Structure of Solids Types of solids: – Crystalline – a well defined arrangement of atoms; this arrangement is often seen on a macroscopic level. Ionic solids – ionic bonds hold the solids in a regular three dimensional arrangement. Molecular solid – solids like ice that are held together by intermolecular forces. Covalent network – a solid consists of atoms held together in large networks or chains by covalent networks. Metallic – similar to covalent network except with metals. Provides high conductivity. – Amorphous – atoms are randomly arranged. No order exists in the solid.