LIQUIDS & SOLIDS

1. LIQUIDS & SOLIDS CHAPTER 10

2. INTERMOLECULAR FORCES • The forces with which molecules attract each other. • Intermolecular forces are weaker than ionic or covalent bonds. • Intermolecular forces are responsible for the physical state of a compound (solid, liquid or gas). • TYPES OF INTERMOLECULAR FORCES • Dipole Interactions (between polar molecules) • London Dispersion Forces (between all molecules but mainly force between nonpolar molecules and noble gases) • Hydrogen Bonds (between molecules where hydrogen is bonded to nitrogen, oxygen and fluorine)

3. DIPOLE-DIPOLE FORCES • Dipole-dipole forces exist between neutral polar molecules. • Polar molecules need to be close together. • Weaker than ion-dipole forces. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

4. LONDON DISPERSION FORCES • Weakest of all intermolecular forces. • It is possible for two adjacent neutral molecules to affect each other. • The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).For an instant, the electron clouds become distorted.In that instant a dipole is formed (called an instantaneous dipole). • Polarizability is the ease with which an electron cloud can be deformed. • The larger the molecule (the greater the number of electrons) the more polarizable. • London dispersion forces increase as molecular weight increases. • London dispersion forces exist between all molecules. • London dispersion forces depend on the shape of the molecule. • The greater the surface area available for contact, the greater the dispersion forces.

5. HYDROGEN BONDING • H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). • Electrons in the H-X (X = electronegative element) lie much closer to X than H. • H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X. • Therefore, H-bonds are strong. • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong.

6. Hydrogen bonds are responsible for: • Ice Floating • Solids are usually more closely packed than liquids; • Therefore, solids are more dense than liquids. • Ice is ordered with an open structure to optimize H-bonding. • Therefore, ice is less dense than water. • In water the H-O bond length is 1.0 Å. • The O…H hydrogen bond length is 1.8 Å. • Ice has waters arranged in an open, regular hexagon. • Each + H points towards a lone pair on O.

7. EFFECTS OF INTERMOLECULAR FORCES ON PHYSICAL PROPERTIES

8. VISCOSITY & SURFACE TENSION • Viscosity • Viscosity is the resistance of a liquid to flow. • A liquid flows by sliding molecules over each other. • The stronger the intermolecular forces, the higher the viscosity. • Surface Tension • Bulk molecules (those in the liquid) are equally attracted to their neighbors.

9. Surface molecules are only attracted inwards towards the bulk molecules. • Therefore, surface molecules are packed more closely than bulk molecules. • Surface tension is the amount of energy required to increase the surface area of a liquid. • Cohesive forcesbind molecules to each other. • Adhesive forcesbind molecules to a surface.

11. BOILING POINT • The stronger the intermolecular forces, the higher the boiling point. • For compounds with approximately the same molecular weight:

12. HYDROGEN BONDED TO GROUP 16 ELEMENTS: NOTICE THAT H2O HAS A GREATER BOILING POINT THAN THE OTHER EVEN THOUGH IT HAS THE LOWEST MOLECULAR MASS DUE TO THE HYDROGEN BONDING. THE OTHER HYDROGEN COMPOUNDS EXPERIENCE DIPOLE-DIPOLE BONDING AND THE BOILING POINT INCREASES WITH INCREASING MOLECULAR MASS. BOILING POINT TRENDS HYDROGEN BONDED TO GROUP 14 ELEMENTS: NOTICE THAT THESE MOLECULES WITH HYDROGEN ARE EXPERIENCE LONDON DISPERSION FORCES AND ALL BOILING POINTS INCREASE WITH INCREASE MOLECULAR MASS

15. A. Evaporation and Vapor Pressure • Vaporization or evaporation • Endothermic

16. EVAPORATION MOLECULES AT THE SURFACE HAVE LESS INTERMOLECULAR FORCES ON THEM THAN THE MOLECULES BELOW THEM . THUS THEY CAN ESCAPE THE LIQUID PHASE EASIER

17. Explaining Vapor Pressure on the Molecular Level • Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid. • These molecules move into the gas phase. • As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. • After some time the pressure of the gas will be constant at the vapor pressure.

18. A. Evaporation and Vapor Pressure • Vapor Pressure • Amount of liquid first decreases then becomes constant • Condensation - process by which vapor molecules convert to a liquid • When no further change is visible the opposing processes balance each other - equilibrium

19. A. Evaporation and Vapor Pressure • Vapor Pressure • Vapor pressure - pressure of the vapor present at equilibrium with its liquid • Vapor pressures vary widely - relates to intermolecular forces

20. Boiling Point and Atmospheric Pressure • BOILING OCCURS WHEN THE VAPOR PRESSURE OF THE LIQUID EQUALS THE ATMOSPHERIC PRESSURE. • LESS PRESSURE MEANS THAT THE MOLECULES CAN ESCAPE EASIER AND THUS HAS A LOWER BOILING POINT

21. Vapor Pressure and Boiling Point • Liquids boil when the external pressure equals the vapor pressure. • Temperature of boiling point increases as pressure increases. • Two ways to get a liquid to boil: increase temperature or decrease pressure. • Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. • Normal boiling point is the boiling point at 760 mmHg (1 atm).

23. MELTING POINT • The melting point is the temperature at which a solid is converted to its liquid phase. • In melting, energy is needed to overcome the attractive forces in the more ordered crystalline solid. • The stronger the intermolecular forces, the higher the melting point. • Because ionic compounds are held together by extremely strong interactions, they have very high melting points. • With covalent molecules, the melting point depends upon the identity of the intermolecular force. For compounds of approximately the same molecular weight:

24. C. Energy Requirements for the Changes of State • Molar heat of fusion – energy required to melt 1 mol of a substance • Molar heat of vaporization – energy required to change 1 mol of a liquid to its vapor

25. Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: • it takes more energy to completely separate molecules, than partially separate them.

26. CALCULATING HEAT OF VAPORIZATION & VAPOR PRESSURE OF WATER Ln(Pvap) =[ (-ΔHvap/R)(1/T)] Textbook : pg 486

27. TO FIND THE BOILING POINT AT A DIFFERENT PRESSURE Ln (P1/P2) = (ΔH/R)[ (1/T2) –(1/T1)] or Ln (P1/P2) = (ΔH/R)[ (1/T2) –(1/T1)] Ln (P2/P1) = (ΔH/R)[ (1/T1) –(1/T2)] ΔH is ΔHvap

29. HEATING AND COOLING CURVES

30. Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions. • The sequence • heat solid  melt  heat liquid  boil  heat gas • is endothermic. • The sequence • cool gas  condense  cool liquid  freeze  cool solid • is exothermic.

31. Plot of temperature change versus heat added is a heating curve. • During a phase change, adding heat causes no temperature change. • These points are used to calculate Hfus and Hvap. • Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. • Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

34. PHASE DIAGRAMS

35. PHASE DIAGRAMS • Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases. • Given a temperature and pressure, phase diagrams tell us which phase will exist. • Any temperature and pressure combination not on a curve represents a single phase.

36. Features of a phase diagram: • Triple point: temperature and pressure at which all three phases are in equilibrium. • Vapor-pressure curve: generally as pressure increases, temperature increases. • Critical point: critical temperature and pressure for the gas. • Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. • Normal melting point: melting point at 1 atm.

37. Critical Temperature and Pressure • Gases liquefied by increasing pressure at some temperature. • Critical temperature: the minimum temperature for liquefaction of a gas using pressure. • Critical pressure: pressure required for liquefaction.

39. Water: • The melting point curve slopes to the left because ice is less dense than water. • Triple point occurs at 0.0098C and 4.58 mmHg. • Normal melting (freezing) point is 0C. • Normal boiling point is 100C. • Critical point is 374C and 218 atm. • Carbon Dioxide: • Triple point occurs at -56.4C and 5.11 atm. • Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.) • Critical point occurs at 31.1C and 73 atm.